Essentials of Inorganic Chemistry ( PDFDrive )

32 Essentials of Inorganic Chemistry

H He

Li Be B C N O F Ne
Na Mg AI Si P S CI Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La- Hf Ta W Re Os Ir Pt Au Hg TI Pb Bi Po At Rn
Lu

Fr Ra Ac- Rf Db Sg Bh Hs Mt Ds Rg Uub
Lr

Figure 2.8 Atomic radii

H He

Li Be B C N O F Ne

Na Mg AI Si P S CI Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La- Hf Ta W Re Os Ir Pt Au Hg TI Pb Bi Po At Rn
Lu

Fr Ra Ac- Rf Db Sg Bh Hs Mt Ds Rg Uub
Lr

Electronegativity

Figure 2.9 Electronegativity

elements diagonally positioned within the periodic table have similar properties, such as similar atomic size,

electronegativity and ionisation energy (Figure 2.11).

The concept of the diagonal relationship is crucial for the biological activity of lithium drugs, which is
mainly due to the properties of the Li+ ion being similar to the Mg2+ ion. In comparison, the size of the Li+
ion is similar to that of Mg2+ and therefore they compete for the same binding sites in proteins. Nevertheless,
lithium has relatively specific effects, and so only proteins with a low affinity for Mg2+ are targeted. Li+
and Mg2+ salts have similar solubility, for example, CO32−, PO43−, F− salts have a low water solubility, and
halide and alkyl salts are soluble in organic solvents. Li+ and Mg2+ compounds are generally hydrated, for

example, LiCl⋅3H2O and MgCl2⋅6H2O. Similarities in ionic size, solubility, electronegativity and solubility
result in similar biological activity and therefore pharmaceutical application [3b].

Alkali Metals 33

H Electronegativity/ He
Li Be ionisation energy

B C N O F Ne

Na Mg AI Si P S CI Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

La-
Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg TI Pb Bi Po At Rn

Fr Ra Ac- Rf Db Sg Bh Hs Mt Ds Rg Uub
Lr

Atomic radii
Figure 2.10 Periodicity showing the ‘metallic character’ trend (highlighted in grey) within the periodic table

H He

Li Be B C N O F Ne

Na Mg AI Si P S CI Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La- Hf Ta W Re Os Ir Pt Au Hg TI Pb Bi Po At Rn
Lu

Fr Ra Ac- Rf Db Sg Bh Hs Mt Ds Rg Uub
Lr

Figure 2.11 Diagonal relationship

2.2.5 What are the pharmacological targets of lithium?

The precise mechanism of action of lithium ions as mood stabilisers is unknown. Current research shows that
there are two main targets in the cell – the enzymes glycogen synthase kinase-3 (GSK-3) and the phospho-
monoesterases family (PMEs) [3b].

GSK-3 is a serine/threonine protein kinase, which is known to play an important role in many biological
processes. GSK-3 is an enzyme that mediates the addition of phosphate molecules onto the hydroxyl groups
of certain serine and threonine amino acids, in particular cellular substrates. Li+ inhibits GSK-3 enzymes via
competition for Mg2+ binding [8].

Phosphoric monoester hydrolases (PMEs) are enzymes that catalyse the hydrolysis of O—P bonds by
nucleophilic attack of phosphorous by cysteine residues or coordinated metal ions. Inositol monophosphatase
(InsP) is the best known member of this family. Li+ and Mg2+ both have a high affinity to bind to phos-
phate groups. Li+ inhibits the enzymatic function of InsP and prevents phosphate release from the active

34 Essentials of Inorganic Chemistry

site. Generally, inositol phosphatases are Mg2+-dependent, but Li+ binds to one of the catalytic Mg2+ sites.
Binding Li+ to phosphate-containing messenger molecules could perturb the transcellular communication
and thus be antipsychotic [3b].

Lithium inhibits GSK-3 and InsP, and both pathways have therefore been suggested to be involved in the
treatment of BD and schizophrenia. The theory behind this hypothesis is that overactive InsP signalling in the
brain of these patients potentially causes BD and this may be reduced by the inhibitory effect of lithium on
such signalling [9].

GSK-3 alters structure of cerebella neurons, and lithium is a neuroprotective agent that can reduce the
hypersensitivity to toxins as seen in cells that overexpress GSK-3. It is believed that lithium potentially can
protect against disease-induced cell death. GSK-3 has been implicated in the origins of schizophrenia, but
with the availability of many antipsychotic drugs on the market, lithium ions are not in common use for the
treatment of schizophrenia [10].

There are also several direct roles of lithium in the treatment of Alzheimer’s disease. Alzheimer’s disease
is a neurodegenerative brain disorder causing neuronal dysfunction and ultimately cell death. This leads ulti-
mately to dementia, affecting in the United States around 10% of people aged over 65 and 48% aged over 85.
Onset occurs with the accumulation of extracellular senile plaques composed of amyloid-β peptides and with
the accumulation of intercellular neurofibrillary tangles. GSK-3 is necessary for the accumulation of tangles,
and subsequently GSK-3 inhibition reduces the production of peptides – this is where the Li+ interaction
comes into play [3b].

2.2.6 Adverse effects and toxicity

Lithium has a very narrow therapeutic window, which makes the monitoring of blood levels essential during
the treatment (see Section 2.2.3). Blood plasma concentrations of more than 1.5 mmol/l Li+ may cause toxic
effects, usually tremors in the fingers, renal impairment and convulsion. Also, memory problems are a very
common side effect, and complaints about slowed mental ability and forgetfulness are commonly reported.
Extreme doses of lithium can cause nausea and diarrhea, and doses above 2 mmol/l require emergency treat-
ment, as these levels may be fatal. Weight gain and decreased thyroid levels are also commonly reported
problems. The blood serum level of Li+ has to be monitored 12 h after administration, and health care pro-
fessionals also have to be aware that the mood-stabilising effects of lithium take a couple of days to take
effect [6].

Lithium salts have severe adverse effects on the renal system. Lithium therapy can damage the internal
structures of the kidneys, which are the tubular structure of the nephrons, and can lead to diabetes insipidus.
One out of five patients experience polyuria – excess production of urine, which is the number one symptom
of diabetes insipidus. Therefore, the kidney function has to be closely monitored for patients undergoing
long-term therapy, with a full test of the kidney function every 6–12 months for stabilised regimes [3a, 6].

Great care has to be taken when lithium is given with nonsteroidal anti-inflammatory drugs (NSAIDs)
because up to 60% increase of Li+ concentration in blood has been observed. NSAIDs reduce the clearance
of lithium through kidneys, and as a result lithium poisoning is possible. Furthermore, the concurrent use of
diuretics, which can result in sodium depletion, can make lithium toxicity worse and can be hazardous [3a].

2.3 Sodium: an essential ion in the human body

Sodium has atomic number 11 and has the symbol Na, derived from the Latin name ‘natrium’. Sodium ions
(Na+) are soluble in water and therefore present in large quantities in the oceans. Na+ is also part of minerals
and an essential element for all animal life.

Alkali Metals 35

The main biological roles of sodium ions are the maintenance of body fluids in humans and the functioning
of neurons and transmission of nerve impulses. Na+ is an important electrolyte and a vital component of the
extracellular fluid. Therefore, one of its roles is to maintain the fluid in the human body via osmoregulation,
a passive transport mechanism (Section 2.3.1). Na+ ions also play a crucial role in the contraction of muscles
and in the mode of action of several enzymes. In the human body, Na+ is often used to actively build up an
electrostatic potential across membranes, with potassium ions (K+) being the counter-ion (Section 2.3.2).
The build-up of an electrostatic potential across cell membranes is important to allow the transmission of
nerve impulses.

2.3.1 Osmosis

Osmosis is defined as the physical process of diffusion of a solvent (water) through a semi-permeable
membrane towards an area of high solute (salt) concentration. This means that solvent (water) follows
the osmotic gradient by moving across the semi-permeable membrane from one solution where there is a
lower salt concentration towards a second solution with a high salt concentration in order to dilute this
and to equalise the concentrations.

Sodium is an essential mineral for the human body and crucial for the regulation of the body fluid via
its osmosis activity. Sodium ions account for over 90% of all ions in the plasma and in the interstitial fluid,
which are involved in osmosis processes. Furthermore, it is the most abundant cation in the extracellular fluid,
and therefore the Na+ content controls the extracellular volume. In particular, the kidneys play an important
role in regulating the fluid level of the body as well as the filtration, secretion and re-absorption of Na+ in
the nephrons, the functional unit of the kidney. Na+ ions are used in the human body to establish osmotic
gradients, which in turn is crucial to control the water balance. Furthermore, decreases in blood pressure and
in Na+ concentrations are sensed by the kidneys, and hormones (e.g. renin, antidiuretic hormones (ADHs),
atrial natriuretic peptide) are released that control the blood pressure, osmotic balances and water-retaining
mechanisms.

In general, if a medium is

• hypertonic, that means the solution has a higher concentration of solutes than the surrounding area. This
area will lose water through osmosis;

• isotonic, that means the solution has the same concentration of solutes as the surrounding area. No move-
ment of water will occur;

• hypotonic, that means the solution has a lower concentration of solutes than the surrounding area. This
area will gain water through osmosis (Figure 2.12).

A net movement of the solvent (water) occurs from the hypotonic solution to the solution with the higher
concentration in order to reduce the difference in concentrations. The osmotic pressure is defined as the
pressure that is required to establish equilibrium with no movement of solvents. It is important to mention
that osmotic pressure depends on the number of ions or molecules in the solution, not the identity of those.
The unit often being found to describe the osmotic pressure is the osmole (osmol or osm), which is a non-SI
unit that defines the numbers of moles of a compound that contributes to the osmotic pressure of a solution.

In general, osmotic processes are important for many biological processes. Plants use osmosis to transport
water and solutes through their systems and the osmotic gradient to establish the turgor within cells. The
human body uses osmosis for many processes, the excretion of urine being one of the most prominent one.

36 Essentials of Inorganic Chemistry

Semipermeable
membrane

Osmosis

Concentrated sugar Diluted sugar

solution solution

Water molecule
Sugar molecule

Figure 2.12 Schematic representation of osmosis

Proximal convoluted tubule Renal capsule
Peritubular capillary Renal corpuscle:

Glomerular (Bowman’s)
capsule
Glomerulus

Efferent arteriole

Distal convoluted tubule

Afferent arteriole
Interlobular artery

Interlobular vein

Renal cortex Renal cortex Arcuate vein
Renal medulla Renal medulla Arcuate artery
Renal papilla Corticomedullary junction

Minor calyx Loop of Henle:
Descending limb
Ascending limb

Collecting duct

Kidney

Figure 2.13 The kidney and its functional unit – the nephron [11] (Reproduced with permission from [11]. Copy-
right © 2009, John Wiley & Sons, Ltd.)

Alkali Metals 37

Urine production takes place within the kidney, more specifically at the nephrons which are the functional
units of the kidneys. Approximately 150–180 of plasma is filtered every day through the glomerulus, which is
a part of the nephron, in order to produce the urine. The nephron also consists of the proximal tubule, the Loop
of Henle and the distal tubule, which leads to the collecting duct and ultimately to the ureter (Figure 2.13) [11].

Filtration takes places at the glomerulus, whereas the remaining parts of the nephron are responsible for the
secretion and re-absorption of ions in order to regulate imbalances and manage the urine volume before the
urine is stored in the bladder. This secretion and re-absorption can occur via an active or a passive transport
across the nephron membrane. Na+ is usually actively transported across via Na+ pumps in order to establish
the correct Na+ concentration in the blood plasma, which is responsible for maintaining the correct osmotic
pressure. Via this process, an osmotic gradient is established within the kidney parenchyma, which is used to
conserve water. The ascending limb of the Loop of Henle is impermeable to water but permeable to Na+. As
a result, an osmotic gradient is established. The descending limb of the Loop of Henle is permeable to water
and, as a result of the osmotic gradient, water moves to the interstitial fluid and urine is concentrated. The
collecting ducts can be permeable to water if the body sends out a signal that water has to be conserved. Again,
water will passively follow the osmotic gradient and urine will be concentrated even more (Figure 2.14).

2.3.2 Active transport of sodium ions

As previously mentioned, the active transport of sodium ions is crucial for the functioning of, for example,
neurons and the subsequent transmission of a nerve impulse. This can be achieved by the active build-up
of a concentration gradient along the cell membrane using Na+/K+ pumps as the active unit. This active

Proximal tubule Distal tubule
HCO3– H2ONaCl K+
H2O
H+ HCO3–

Glomerulus Collecting duct 300
600
H+ NH3 NaCl Na+ 900
NaCl 1200

H2O Loop of Henle NaCl
H2O NaCl
Passive transport H2O NaCl
Active transport H2O NaCl

Interstitial fluid
osmolality

Figure 2.14 Osmotic gradient in kidney parenchyma

38 Essentials of Inorganic Chemistry

Na+ Extracellular fluid 3Na+ expelled
gradient
Na+/K+ ATPase 2K+

Cytosol 3Na+ P 3P 4 2K+
K+ 1 ATP imported
gradient
2
ADP

Figure 2.15 Mode of action of the Na+/K+-ATPase [11] (Reproduced with permission from [11]. Copyright ©
2009, John Wiley & Sons, Ltd.)

transport is responsible for the cells containing relatively high concentrations of potassium ions and low

concentrations of sodium ions. The resulting electrostatic potential that is built up along the cell membrane

is called action potential and is subsequently responsible for the transmission of nerve impulses (see Section

2.4.1) (Figure 2.15).
The Na+/K+ pumps facilitate an active transport process which is based on the conformational changes of

the cross-membrane protein and driven by the breakdown of ATP. In the initial step, three Na+ ions bind the

cross-membrane protein on the cytosolic side. This causes the protein to change its confirmation and makes

it accessible to ATP. In its new confirmation, the protein becomes phosphorylated by ATP, which results in
a second conformational change. The three Na+ ions are located across the membrane, and the protein now

has a low affinity to the sodium ions. This means that the sodium ions are dissociated from the protein and
released into the extracellular fluid. Nevertheless, the protein has now a high affinity to K+ and binds two

potassium ions from the extracellular fluid. The bond phosphate is now dissociated, and the protein reverts
back to its original confirmation. This means both K+ ions are exposed to the cytosol and can be released.

2.3.3 Drugs, diet and toxicity

Sodium chloride solutions are normally used when the patient is diagnosed with sodium depletion and dehy-
dration. Treatment is mostly administered intravenously, but in chronic conditions (mild to moderate sodium
loss) sodium chloride or sodium bicarbonate can be given orally. Oral rehydration therapies usually use a
mixture of alkali metal-based salts such as NaCl, KCl and their citrates (Figure 2.16) [6].

Sodium bicarbonate is usually administered orally in order to regulate the serum pH. Imbalances of the
plasma pH can be due to problems occurring in the kidneys such as renal tubular acidosis. This is a medical
condition that occurs where the body accumulates acid as a result of the kidneys failing to regulate the pH of
the urine and the blood plasma. Within the kidneys, blood is filtered before it passes through the tubular part
of the nephrons where re-absorption or secretion of important salts and others takes place. In renal tubular
acidosis, the kidneys either fail to filter or secrete acid ions (H+) from the plasma (secretion takes place in
the distal tubule), or to recover bicarbonate ions (HCO3−) from the filtrate (passive re-absorption takes place
in the proximal tubule, active re-absorption at the distal tubule), which is necessary to balance the pH. In the
view of this mode of action, the pharmaceutically active component of sodium bicarbonate is the bicarbonate
anion, but the cation Na+ is responsible for solubility and compatibility (Figure 2.17) [3a].

The most common dietary source of NaCl is table salt, which is used for seasoning and pickling (the
high NaCl content inhibits the bacterial and fungal growth as a result of the osmotic gradient). The daily
recommended NaCl intake varies depending on the country and the age group. Within the United Kingdom,

Alkali Metals 39

NaCl O O OH O
Na+ Na+ Na+
O
O OH O

(a) (b) OO
Na+

(c)

Figure 2.16 Chemical structures of (a) sodium chloride, (b) sodium bicarbonate and (c) sodium citrate

Proximal tubule Distal tubule
HCO3– H2ONaCl K+
H2O
Glomerulus H+ HCO3– Collecting duct 300
600
K+ 900
1200
H+ NH3 NaCl Na+
NaCl

H2O Loop of Henle NaCl
H2O NaCl
Passive transport H2O NaCl
Active transport H2O NaCl

Interstitial fluid
osmolality

Figure 2.17 Illustration of a nephron showing areas of active and passive electrolyte transport

the maximum salt intake is recommended to be limited to 6 g of NaCl for an adult, whereas intake for children
should be significantly lower [12]. Most people exceed this amount on a daily basis, and the high salt plasma
levels (hypernatraemia) can result in cardiovascular disorders such as hypertension. Low sodium plasma levels
(hyponatraemia), which again can be a result of dysfunction kidneys or sodium loss in the bowels, also cause
damage to the human body via osmotic imbalances and if necessary have to be treated. Low blood pressure,
dehydration and muscle cramps are signs of a sodium deficiency.

Signs of acute toxicity may be seen after ingestion of 500–1000 mg/kg body weight NaCl. These symptoms
can be vomiting, ulceration of the gastrointestinal (GI) tract and renal damage. Also, the increased risk for
the formation of kidney stones is believed to be a result of high salt intake [13].

40 Essentials of Inorganic Chemistry

2.4 Potassium and its clinical application

Potassium has atomic number 19 and the chemical symbol K, which is derived from its Latin name ‘kalium’.
Potassium was first isolated from potash, which is potassium carbonate (K2CO3). Potassium occurs in nature
only in the form of its ion (K+) either dissolved in the ocean or coordinated in minerals because elemental
potassium reacts violently with water (see Section 2.1.3). Potassium ions are essential for the human body
and are also present in plants. The major use of K+ can be found in fertilisers, which contains a variety of
potassium salts such as potassium chloride (KCl), potassium sulfate (K2SO4) and potassium nitrate (KNO3).
KCl is also found in table salt, whereas potassium bromate (KBrO3) is an oxidising agent and is used as flour
improver. Potassium bisulfite (KHSO3) can be used as a food preservative in wine and beer.

2.4.1 Biological importance of potassium ions in the human body – action potential

The so-called action potential occurs in a variety of excitable cells such as neurons, muscle cells and endocrine
cells. It is a short-lasting change of the membrane potential and plays a vital role in the cell-to-cell communi-
cation. In animal cells, there are two types of action potential. One type is produced by the opening of calcium
ion (Ca2+) channels and this is longer lasting than the second type, namely the Na+-based action potential.

The Na+/K+-based action potential is short-lived (only 1 ms) and therefore mostly found in the brain and
nerve cells. Potassium ions are crucial for the functioning of neurons, by influencing the osmotic balance
between the cells and the interstitial fluid. The concentration of K+ within and outside the cells is regulated
by the so-called Na+/K+-ATPase pump. Under the use of ATP, three Na+ ions are pumped outside the cell and
two K+ ions are actively transported into the cell (see Section 2.3.2). As a result, an electrochemical gradient
over the cell membrane is created; the so-called resting potential is established.

In the case of any cell-to-cell communication, changes of the membrane reach a specific part of the neu-
ron first. As a result, sodium channels, which are located in the cell membrane, will open. As a result of
the osmotic gradient, which has been established by the Na+/K+-pump, Na+ ions enter the cytosol and the
electrochemical gradient becomes less negative. Once a certain level, called the threshold, is reached, more
sodium channels are opened, and more Na+ ions flow inside the cell – this creates the so-called action poten-
tial. The electrochemical gradient over the cell membrane is thus reversed. Once this happens, the sodium
ion channels close and the potassium channels open. The concentration of K+ in the cytosol is higher than
in the extracellular fluid, and therefore potassium ions leave the cell. This allows the cell to shift back to its
resting membrane potential. As the membrane potential approaches the resting potential, all voltage-gated
K+ channels open. In actual fact, the membrane repolarises beyond the resting potential; this is known as
hyperpolarisation. The last step is now to reintroduce the initial balance of sodium and potassium ions. This
means that the Na+/K+-pump transports Na+ actively out of the cytosol and K+ into it. As a result, the initial
steady state is reinstated (Figure 2.18).

2.4.2 Excursus: the Nernst equation

As previously outlined, the electric potential across a cell membrane is created by the difference in ion concen-
tration inside and the outside the cell. The Nernst equation is an important equation that allows the calculation
of the electric potential for an individual ion.

ΔE = − RT ln [ion]in
zF [ion]out

Alkali Metals 41

3

Membrane potential (mV) 30 4
5
0
2

–30
1

6

Figure 2.18 Action potential. (1) Threshold of excitation. Na+ channels open and allow Na+ to enter the cell.
(2) K+ channels open and K+ leaves the cell. (3) No more Na+ enters cell. (4) K+ continues to leave the cell.
This causes the membrane to return to resting potential. (5) K+ channels close and Na+ channels resent. (6)

Hyperpolarisation

Looking at the Nernst equation, R is the gas constant, T is the temperature in kelvin, F is the Faraday constant

and z is the net charge of the ion. [ion]in and [ion]out are the concentration of the particular ion inside and
outside the cell. The equation can also be expressed as log10. Furthermore, at room temperature (25 ∘C), the

term −RT/F can be seen as a constant, and the Nernst equation can be simplified:

ΔE = − 0.059 log [ion]in
zV [ion]out

The Nernst equation can also be used to calculate the overall potential in a redox equation by using the ion
concentrations of both half-equations. This also allows the calculation of the reduction potential ΔE of two
different ions of varying concentration. The Nernst equation can be expanded to the following equation, where
[Red] and [Ox] represent the concentrations of the reductant and oxidant (Ox + e− → Red):

ΔE = E0 − 0.059 log [Red]
zV [Ox]

In order to calculate the reduction potential ΔE of two different ions of varying concentration, the Nernst
equation can be expanded to the following equation:

ΔE = ΔE0 − 0.059 log [ion]in
zV [ion]out

It is also possible to calculate the overall electric potential across the cell membrane by taking several
different ions into account. The so-called Goldman equation, which will not be further illustrated in this
book, can be used to calculate this value.

42 Essentials of Inorganic Chemistry

2.4.3 Potassium salts and their clinical application: hypokalaemia

In the human body, 95% of the K+ can be found inside the cells, with the remaining 5% mainly circulating
in the blood plasma [11]. This balance is carefully maintained by the Na+/K+ pump (see Section 2.3.2), and
imbalances, such as seen in hypo or hyperkalaemia, can have serious consequences.

Hypokalaemia is a potentially serious condition where the patient has low levels of K+ in his/her blood
plasma. Symptoms can include weakness of the muscles or ECG (electrocardiogram) abnormalities. Mostly,
hypokalaemia can be a result of reduced K+ intake caused by GI disturbance, such as diarrhoea and vomit-
ing, or increased excretion of K+ caused by diuresis. Hypokalaemia is often found in patients treated with
diuretics such as loop diuretics and thiazides. These classes of drugs increase the secretion of Na+ in the
nephrons in order to increase water excretion. Unfortunately, they also increase the excretion of K+ and
lead to hypokalaemia. In contrast, potassium-sparing diuretics actively preserve potassium ions, and patients
treated with loop diuretics or thiazides often receive also potassium-sparing diuretics [3a].

Potassium ions are excreted via the kidneys. Within the kidneys, ∼150–180 l of plasma is filtered every day
through the glomerulus, which is part of the nephron, in order to produce urine. As previously described, the
filtration process is followed by a series of processes along the nephron, where a variety of ions are secreted
and re-absorbed in order to regulate plasma imbalances and manage the urine volume (Section 2.3.1). K+ is
passively secreted at the proximal tubule and also moves into the interstitial fluid via a counter-flow process
to Na+ mainly at the distal tubule (Figure 2.19).

Oral supplementation in form of potassium salts is especially necessary in patients who take anti-arrhythmic
drugs, suffer from renal artery stenosis and/or severe heart failure or show severe K+ losses due to chronic

Proximal tubule Distal tubule
HCO3– H2ONaCl K+
H2O
Glomerulus H+ HCO3– Collecting duct 300
600
K+ 900
1200
H+ NH3 NaCl Na+
NaCl

H2O Loop of Henle NaCl
H2O NaCl
Passive transport H2O NaCl
Active transport H2O NaCl

Interstitial fluid
osmolality

Figure 2.19 Illustration of a nephron showing areas of potassium transport

Alkali Metals 43

O
K+

O OH

Figure 2.20 Chemical structure of potassium bicarbonate (KCO3H)

diarrhoea or abusive use of laxatives. Regulation of the plasma K+ level may also be required in the care of
elderly patients when the K+ intake is reduced as a result of changing dietary habits, but special attention has
to be given to patients with renal insufficiency because K+ excretion might be reduced. Potassium salts are
preferably given as liquid preparations, and KCl is the preferred salt used. Other potassium-based salts can
be used if the patient is at risk of developing hyperchloraemia – increased chloride plasma levels. Typically
potassium salts are dissolved in water, but the salty and bitter taste makes them difficult to formulate. Oral
bicarbonate solutions such as potassium bicarbonate are typically given orally for chronic acidosis states – low
pH of the blood plasma. This can be again due to impaired kidney function. The use of potassium bicarbonate
for the treatment of acidosis has to be carefully evaluated, as even small changes of the potassium plasma
levels can have severe consequences (Figure 2.20) [3a].

Potassium citrate is used in the United Kingdom as an over-the-counter drug for the relief from discomfort
experienced in mild urinary-tract infections by increasing the urinary pH. It should be not given to men if
they experience pain in the kidney area (risk of kidney stones) or if blood or pus is present in the urine. Also,
patients with raised blood pressure or diabetes should avoid taking potassium citrate without consultation
with their general practitioner (GP). Caution is generally advised to patients with renal impairment, cardiac
problems and the elderly [3a].

2.4.4 Adverse effects and toxicity: hyperkalaemia

The therapeutic window for K+ in the blood plasma is very small (3.5–5.0 mmol), and especially hyper-
kalaemia, an increased level of K+ in the plasma, can lead to severe health problems [11]. Potassium salts
can cause nausea and vomiting and in extreme cases can lead to small bowel ulcerations. Acute severe hyper-
kalaemia is defined when the plasma potassium concentration exceeds 6.5 mmol/l or if ECG changes are
seen. This can lead to cardiac arrest, which needs immediate treatment. Treatment options include the use
of calcium gluconate intravenous injections, which minimises the effects of hyperkalaemia on the heart. The
intravenous injection of soluble insulin promotes the shift of potassium ions into the cells. Diuretics can also
be used to increase the secretion of K+ in the kidneys, and dialysis can be a good option if urgent treat-
ment is required. Ion-exchange resins, such as polystyrene sulfonate resins, may be used in mild to moderate
hyperkalaemia to remove excess potassium if there are no ECG changes present. As previously mentioned,
especially in patients suffering from kidney diseases or end-stage renal failure, the potassium levels have to
be monitored very carefully and corrected if necessary. Potassium excretion is likely to be disturbed, and a
build-up of potassium in the blood plasma may trigger a cardiac arrest [3a].

Potassium salts are also available in the form of tablets or capsules for oral application especially as nonpre-
scription medicine. Usually, their formulation is designed to allow the potassium ions to be slowly secreted,
because very high concentrations of K+ are known to be toxic to tissue cells and can cause injury to the gastric
mucosa. Therefore, nonprescription potassium supplement pills are usually restricted to <100 mg of K+ [6].



Alkali Metals 45

2.5 Exercises

2.5.1 Determine the oxidation state for all elements in the following molecules:
2.5.2
2.5.3 (a) H2O
(b) NaCl
(c) H2O2
(d) Fe2O3
(e) MnO4−
(f) Cr2O72−
(g) SiH4

Complete the reduction and oxidation equation and write the redox equation in the following
examples:
(a) MnO4− + Fe2++ ???? → Mn2+ + Fe3++ ?????? under acidic conditions

(b) MnO4− + Br− + OH− → ??? + BrO3− under basic conditions

(c) Cr2O72− + Cu+ → Cu2++ ??? under acidic conditions

(d) IO3− + I− → I2 under acidic conditions

Complete the equation and indicate the standard reduction potential E r ed assuming both reac-
tions partners are used in the same concentration.

(a) Br2 + 2I− → ???? E0(Br2∕Br−) = 1.07 V
E0(I2∕I−) = 0.54 V

(b) Fe2+ + Ce4+ → ???? E0(Ce4+∕Ce3+) = 1.61 V
E0(Fe3+∕Fe2+) = 0.77 V

(c) Br2 + Fe → ???? E0(Br2∕Br−) = 1.07 V
E0(Fe2+∕Fe) = −0.44 V

2.5.4 Calculate E for the following redox pair when Mn3+= 0.5 M and Mn2+ = 0.01 M [E0(Mn3+/Mn2+)
= 1.51 V] using the Nernst equation.



Alkali Metals 47

2.6 Case studies

2.6.1 Lithium carbonate (Li2CO3) tablets
Your pharmaceutical analysis company has been contacted by an important client and asked to analyse a batch
of formulated Li2CO3 tablets. The description of your brief states that you are supposed to analyse the API
in these tablets following standard quality assurance guidelines.

Typical analysis methods used for quality purposes are based on titration reactions. A certain amount of
powdered Li2CO3 tablets is dissolved in water, and a known amount of HCl is added. The solution is boiled
to remove any CO2. The excess acid is then titrated with NaOH using methyl orange as an indicator [14].
(a) Research the type of titration described. Describe the chemical structure and mode of action of the

indicator.
(b) Formulate the relevant chemical equations.
(c) The package states that each tablet contains 250 mg Li2CO3. For the experiment, 20 tablets are weighed

and powdered (total weight 9.7 g). Powder containing 1 g of Li2CO3 is dissolved in 100 ml water, and
50 ml of 1 M HCl is added. After boiling, the solution is titrated against 1 M NaOH using methyl orange
as the indicator. For each titration, the following volume of NaOH was used:

35.0 ml 35.5 ml 34.5 ml

Calculate the amount of Li2CO3 present in your sample. Express your answer in grams and moles.
(d) Critically discuss your result in context with the stated value for the API.
(e) Research and critically discuss the typically accepted error margins.

2.6.2 Sodium chloride eye drops

Your pharmaceutical analysis company has been contacted by an important client and asked to analyse a batch
of eye drops containing a NaCl solution. The description of your brief states that you are supposed to analyse
the API in these tablets following standard quality assurance guidelines.

Typical analysis methods used for quality purposes are based on titration reactions. A certain volume
of NaCl solution is titrated with silver nitrate (AgNO3). Potassium chromate is used as the appropriate
indicator [14].
(a) Research the type of titration described. Describe the chemical structure and mode of action of the

indicator.
(b) Formulate the relevant chemical equations.
(c) The package states that the eye drops are a 0.9% w/v aqueous solution of NaCl. For the experiment, a

volume containing 0.1 g of NaCl is titrated with a 0.1 M AgNO3 solution. For each titration, the following
volume of AgNO3 is used:

16.9 ml 17.0 ml 17.4 ml

Calculate the amount of NaCl present in your sample. Express your answer in grams and moles.
(d) Critically discuss your result in context with the stated value for the API.
(e) Research the typically accepted error margins.

48 Essentials of Inorganic Chemistry

References

1. J. D. Lee, Concise inorganic chemistry, 5th ed., Chapman & Hall, London, 1996.
2. A. D. McNaught, A. Wilkinson, Compendium of chemical terminology: IUPAC recommendations, 2nd ed./compiled

by Alan D. McNaught, Andrew Wilkinson. ed., Blackwell Science, Oxford, 1997.
3. (a) G. A. McKay, M. R. Walters, J. L. Reid, Lecture notes. Clinical pharmacology and therapeutics, 8th ed.,

Wiley-Blackwell, Chichester, 2010; (b) E. R. Tiekink, M. Gielen, Metallotherapeutic drugs and metal-based
diagnostic agents: the use of metals in medicine, Wiley, Chichester, 2005.
4. J. F. J. Cade, Med. J. Aust. 1975, 1, 684–686.
5. J. Levine, K. N. R. Chengappa, J. S. Brar, S. Gershon, E. Yablonsky, D. Stapf, D. J. Kupfer, Bipolar Disord. 2000,
2, 120–130.
6. Joint Formulary Committee. British National Formulary. 60 ed. London: BMJ Group and Pharmaceutical Press;
2010.
7. N. Farrell, Uses of inorganic chemistry in medicine, Royal Society of Chemistry, Cambridge, 1999.
8. (a) W. C. Drevets, J. L. Price, J. R. Simpson, R. D. Todd, T. Reich, M. Vannier, M. E. Raichle, Nature 1997,
386, 824–827; (b) D. Ongur, W. C. Drevets, J. L. Price, Proc. Natl. Acad. Sci. U.S.A. 1998, 95, 13290–13295;
(c) G. Rajkowska, J. J. Miguel-Hidalgo, J. R. Wei, G. Dilley, S. D. Pittman, H. Y. Meltzer, J. C. Overholser, B. L.
Roth, C. A. Stockmeier, Biol. Psychiatry 1999, 45, 1085–1098; (d) G. J. Moore, J. M. Bebchuk, K. Hasanat, G. Chen,
N. Seraji-Bozorgzad, I. B. Wilds, M. W. Faulk, S. Koch, D. A. Glitz, L. Jolkovsky, H. K. Manji, Biol. Psychiatry
2000, 48, 1–8.
9. M. J. Berridge, C. P. Downes, M. R. Hanley, Cell 1989, 59, 411–419.
10. (a) E. S. Emamian, D. Hall, M. J. Birnbaum, M. Karayiorgou, J. A. Gogos, Nat. Genet. 2004, 36, 131–137; (b)
T. Katsu, H. Ujike, T. Nakano, Y. Tanaka, A. Nomura, K. Nakata, M. Takaki, A. Sakai, N. Uchida, T. Imamura,
S. Kuroda, Neurosci. Lett. 2003, 353, 53–56; (c) J. Z. Yang, T. M. Si, Y. S. Ling, Y. Ruan, Y. H. Han, X. L. Wang,
H. Y. Zhang, Q. M. Kong, X. N. Li, C. Liu, D. R. Zhang, M. Zhou, Y. Q. Yu, S. Z. Liu, L. Shu, D. L. Ma, J. Wei, D.
Zhang, Biol. Psychiatry 2003, 54, 1298–1301.
11. G. J. Tortora, B. Derrickson, Principles of anatomy and physiology, 12th ed., international student/Gerard J. Tortora,
Bryan Derrickson. ed., Wiley [Chichester: John Wiley, distributor], Hoboken, N.J., 2009.
12. NHS, Vol. 2014, NHS, 2013.
13. Minerals, E. G. o. V. a., Vol. 2013, UK Food Standards Agency, 2003.
14. British pharmacopoeia, Published for the General Medical Council by Constable & Co, London.

Further Reading

1. E. Alessio, Bioinorganic medicinal chemistry, Wiley-VCH, Weinheim, 2011.
2. W. Kaim, B. Schwederski, Bioinorganic chemistry: inorganic elements in the chemistry of life: an introduction and

guide, Wiley, Chichester, 1994.
3. H.-B. Kraatz, N. Metzler-Nolte, Concepts and models in bioinorganic chemistry, Wiley-VCH [Chichester: John

Wiley, distributor], Weinheim, 2006.
4. R. M. Roat-Malone, Bioinorganic chemistry: a short course, Wiley, Hoboken, N.J. [Great Britain], 2002.

3

Alkaline Earth Metals

Members of group 2 of the periodic table (second vertical column) are called earth alkaline metals. In this
group are included the following elements: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr),
barium (Ba) and radium (Ra). Radium is a radioactive element and therefore we will not further discuss it in
this chapter (Figure 3.1).

In terms of clinical use, magnesium and calcium are essential ions for the human body and any of their
imbalances should be corrected. Strontium is medically used in radiotherapy, and its application is further
discussed in Chapter 10. Exposure to excess beryllium can lead to the so-called chronic beryllium disease
(CBD), which is discussed later in this chapter. Barium salts are generally highly toxic. Nevertheless, the
so-called barium meal is a well-used oral radio-contrast agent.

3.1 Earth alkaline metal ions

Earth alkaline metals together with the alkali metals form the so-called s-block metals. Earth alkaline metals
have two electrons in their outer shell which is an s-orbital type. The chemistry of the metals is characterised
by the loss of both electrons, which is a result of the relatively low ionisation energy (IE) of both electrons
and the subsequent formation of the stable cation M2+, which has a noble gas configuration (Table 3.1).

Group 2 elements are all silvery-white metals with high reactivity, similar to alkali metals, but less soft
and not as reactive. Earth alkaline metals can be mostly found in the earth’s crust in the form of their cations
displayed in minerals and not as the elemental metal, as these are very reactive. For example, beryllium
principally occurs as beryl (Be3Al2[Si6O18]), which is also known as aquamarine.

Magnesium can be found in rock structures such as magnesite (MgCO3) and dolomite (MgCO3⋅CaCO3),
and is the eighth most abundant element in the earth’s crust. Calcium is the fifth most abundant element and
can be found in minerals such as limestone (CaCO3) and its metamorphs such as chalk and marble.

Earth alkaline metals are harder and have a higher density than sodium and potassium and higher melting
points. This is mostly due to the presence of two valence electrons and the resulting stronger metallic bond.
Atomic and ionic radii increase within the group, and the ionic radii are significantly smaller than the atomic
radii. Again, this is due to the existence of two valence electrons, which are located in the s orbital furthest
from the nucleus. The remaining electrons are attracted even closer to the nucleus as a result of the increased

Essentials of Inorganic Chemistry: For Students of Pharmacy, Pharmaceutical Sciences and Medicinal Chemistry,
First Edition. Katja A. Strohfeldt.
© 2015 John Wiley & Sons, Ltd. Published 2015 by John Wiley & Sons, Ltd.
Companion website: www.wiley.com/go/strohfeldt/essentials

50 Essentials of Inorganic Chemistry

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La- Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Lu

Fr Ra Ac- Rf Db Sg Bh Hs Mt Ds Rg Uub
Lr

Figure 3.1 Periodic table of elements; group 2 elements are highlighted

Table 3.1 First, second and third ionisation
energies (kJ/mol) of group 2 metals [1]

First Second Third

Be 900 1 757 14 847
Mg 738 1 450 7 731
Ca 590 1 145 4 910
Sr 550 1 064 4 207
Ba 503 3 600
965

Source: Reproduced with permission from [8]. Copyright ©
1996, John Wiley & Sons, Ltd.

effective nuclear charge. The IEs of the first two valence electrons are similar and relatively low compared
to the energy needed to remove the third valence electron, which is part of a fully filled quantum shell. As a
result, the dominant oxidation state of earth alkaline metals is +2.

3.1.1 Major uses and extraction

Beryllium is one of the lightest metals and therefore is used in high-speed aircrafts and missiles. Unfor-
tunately, it is highly toxic, and CBD, a scarring of the lung tissue, is often seen in workers from within a
beryllium-contaminated work environment.

Calcium is mostly found in limestone and its related forms, such as chalk, and marble and lime (CaO). Ca2+
ions are essential for living organisms, as is Mg2+. Magnesium is the only earth alkaline metal that is used on
an industrial scale. It is used in ammunition (e.g. tracer bullets and incendiary bombs), as it burns with a very
bright white glow. Magnesium alloyed with aluminium results in a low-density and strong material, which is
used for lightweight vehicles and aeroplanes.

Magnesium is the only group 2 element that is extracted on a large scale. Its main source is seawater, and
the metal is extracted by adding calcium hydroxide. Magnesium hydroxide precipitates, as it is less soluble
in water compared to the calcium compound. Magnesium hydroxide is converted into magnesium chloride

Alkaline Earth Metals 51

Conversion 2HCl + Mg(OH)2 → MgCl2 + 2H2O
Electrolysis: at the cathode:
Mg2+(l) + 2e– → Mg(l)
at the anode:
2Cl–(l) → Cl2(g) + 2e–

Redox: 2Mg2+ + 2Cl– → 2 Mg(l) + Cl2(g)

Figure 3.2 Redox equation for the production of magnesium

Calcination ∶ Dolomite [CaMg(CO3)2] is converted into MgO and CaO
Reduction ∶ 2MgO + 2CaO + FeSi → 2Mg + Ca2SiO4 + Fe + 1450 K

Figure 3.3 Chemical equation for the production of magnesium

(MgCl2), which can be subsequently electrolysed in a Down’s cell (see Section 2.1.1) in order to produce the
pure magnesium metal (Figure 3.2).

Alternatively, there is a second method called the ferrosilicon process or pigeon process. This involves the
reduction of magnesium oxide, which is obtained from dolomite, with an iron–silicon alloy. The raw material
has to be calcined first, which means the removal of water and carbon dioxide, as these would form gaseous
by-products and would reverse the subsequent reduction (Figure 3.3).

3.1.2 Chemical properties

The chemical behaviour of alkaline earth metals is characterised by their strong reducing power, and there-
fore they very easily form bivalent cations (M2+). The elements within group 2 become increasingly more
electropositive on descending within the group.

The metals themselves are coloured from grey (Be, Mg) to silver (Ca, Sr, Ba) and are soft. Beryllium and
magnesium are passivated and therefore kinetically inert to oxygen or water. The metal barium has to be stored
under oil because of its reactivity. Metals such as calcium, strontium and barium react similar to sodium, but
are slightly less reactive: All the metals except beryllium form oxides in air at room temperature once the
reaction is started. The nitride compound is formed in the presence of nitrogen, and magnesium can burn in
carbon dioxide, which means that magnesium fires cannot be extinguished by the use of carbon dioxide fire
extinguishers.

2M + O2 → 2MO 3M + N2 → M3N2
2Mg(s) + CO2(g) → 2MgO(s) + C(s)

The oxides of alkaline earth metals have the general formula MO and are generally basic. Beryllium oxide
(BeO) is formed by the ignition of beryllium metal in an oxygen atmosphere. The resulting solid is colourless
and insoluble in water. Other group 2 oxides (MO) are typically formed by the thermal decomposition of the
corresponding metal carbonate or hydroxide.

MCO3 → MO + CO2

52 Essentials of Inorganic Chemistry

Be(OH)2 + 2(OH)− → [Be(OH)4]2−
Be(OH)2 + H2SO4 → BeSO4 + 2H2O

Figure 3.4 Chemical equations showing the amphoteric nature of beryllium hydroxide

Peroxides are known for magnesium, calcium, strontium and barium but not for beryllium. The radius of
the beryllium cation (Be2+) is not sufficient to accommodate the peroxide anion.

Beryllium reacts with aqueous alkali (NaOH) and forms beryllium hydroxide, which is an amphoteric
hydroxide.

The term amphoteric describes compounds that can act as an acid and a base.

Beryllium hydroxide reacts with a base with the formation of the corresponding beryllium salt. Beryllium

hydroxide can also be reacted with acids such as sulfuric acid and the corresponding salt, beryllium sulfate,
is obtained (Figure 3.4).

Magnesium does not react with aqueous alkali (NaOH). The synthesis of magnesium hydroxide [Mg(OH)2]
is based on a metathesis reaction in which magnesium salts are reacted with sodium or potassium hydroxide.

Mg2+(aq) + 2KOH(aq) → Mg(OH)2(s) + 2K+ (3.1)

Calcium, strontium and barium oxides react exothermically with water to form the corresponding
hydroxides:

CaO(s) + H2O(l) → Ca(OH)2(s)

SrO(s) + H2O(l) → Sr(OH)2(s)

BaO(s) + H2O(l) → Ba(OH)2(s) (3.2)

Magnesium and calcium hydroxides are sparingly soluble in water and the resulting aqueous solutions are

mildly alkaline. In general, group 2 hydroxides, except Be(OH)2, react as bases, and their water solubility
and thermal stability increase within the group (Mg→Ba).

Earth alkaline halides (MCl2) are normally found in their hydrated form. Anhydrous beryllium halides are
covalent, whereas Mg(II), Ca(II), Sr(II), Ba(II) halides are ionic. As a result of the ionic bond, the later halides

have typically high melting points and they are sparingly soluble in water. Additionally, MgCl2, MgBr2, MgI2
are hygroscopic. Anhydrous calcium chloride also has a strong affinity for water and is typically used as a

drying agent.

Hygroscopic compounds are substances that absorb water from the surrounding air but do not become
a liquid.

3.2 Beryllium and chronic beryllium disease

The element beryllium can be found in the mineral beryl [Be3Al2(SiO3)6] and has minor but important tech-
nical applications. Owing to its unique properties, it is used in industrial lightweight systems, for example,

Alkaline Earth Metals 53

turbine rotor blades, automotive parts and electrical contacts. The pure beryllium metal is also used in the
nuclear industry.

Beryllium has an exceptionally small atomic radius, and as a result beryllium fluoride, chloride and oxide
show evidence of covalent bonds in contrast to the other group 2 oxides or halides. Beryllium halides should
be linear if they exhibit the ionic bonding character. This linear form can only be found in the gas phase. In
the solid state, the beryllium centre is three or fourfold coordinated, which can be achieved, for example, by
polymerisation.

Beryllium and its compounds are extremely poisonous and therefore there is only a very limited potential
for their clinical applications. Indeed, even the inhalation of beryllium or its compounds can lead to serious
respiratory diseases such as the chronic beryllium disease, and soluble beryllium compounds can cause serious
skin irritations. Workers within the metal production industry are most likely exposed to beryllium and run
the highest risk of developing CBD. But also people working in connected professions such as administrative
staff or families are at high risk of beryllium poisoning. Symptoms are not well reported, may occur many
years after the exposure and include cough, fatigue and chest pain, whereas nonrespiratory organs can also
be affected. However, the introduction of exposure limits and general awareness of the risk have significantly
reduced the risk of beryllium exposure and its consequences [2].

3.3 Magnesium: competition to lithium?

The element magnesium (Mg) is a silvery-white and lightweight metal. It is protected by a thin oxide layer,
which is very difficult to remove but at the same time removes the need to store it in an oxygen-free envi-
ronment (see alkali metals). Magnesium reacts with water but much more slowly than its neighbouring earth
alkaline metal calcium. Magnesium is a highly flammable metal, and once ignited it burns with a characteris-
tic bright white flame. There are three stable isotopes of magnesium, namely 24Mg (79% occurrence), 25Mg
and 26Mg. 28Mg is radioactive with a half-life of 21 h [1].

Most magnesium salts are soluble in water, and given in large amounts they work as a laxative in the human
body. Aqueous magnesium ions are sour in taste. Magnesium hydroxide (MgOH2) has only limited solubility
in water and the resulting suspension is called milk of magnesia, which is commonly used as an antacid and is
known to be a mild base. Magnesium is extracted on a large scale using a Down’s cell (see Section 2.1.1) or
the ferrosilicon process with seawater being the main source (see Section 3.1.1). Mg2+ stands in a so-called
diagonal relationship to Li+, which explains why these ions have similar properties and biological activity
(see Section 2.2.4).

3.3.1 Biological importance

Mg2+ is an essential ion in the human body and is a crucial constituent in numerous enzymatic processes.
Indeed, Mg2+ is essential to most living cells as a signalling molecule and is involved in nucleic acid bio-
chemistry dealing with the manipulation of ATP (adenosine triphosphate), DNA, RNA and related processes.
For example, ATP has to be coordinated to a magnesium ion in order to become biologically active. Mg2+
also stabilises DNA and RNA structures, which can be seen in their increased melting points.

Mg2+ ions form the redox-active centre in chlorophyll, which facilitates the process of photosynthesis and
the connected carbon fixation in green plants. Therefore, green vegetables, as well as milk, whole grain and
nuts, are good sources of magnesium. It has to be kept in mind that most magnesium salts are water soluble
and therefore processed vegetables, mainly cooked in water, are low in magnesium ion content.

In the human body, Mg2+ is the fourth most abundant cation and the second most abundant ion in the
interstitial fluid. Mg2+ is an essential co-factor dealing with more than 300 cellular enzymatic processes.

54 Essentials of Inorganic Chemistry

On average, the human body contains about 24 g of magnesium ions, with half of it being incorporated into
bones and the other half being present in muscles and soft tissue. The majority of Mg2+ is absorbed in the ilium
and colon, and the kidneys are the major excretory organ. Mg2+ is filtered at the glomerulus, and 10–15% is
re-absorbed at the proximal tubule, 60–70% at the thick part of the ascending limb of the loop of Henle and
10–15% at the distal tubule [3].

Nevertheless, magnesium salts are generally not well absorbed from the gastrointestinal (GI) tract and
therefore are often used as osmotic laxatives. The kidneys regulate the magnesium ion levels in plasma, and
as a result high levels of Mg2+ are retained when the patient has renal failure. The resulting hypermagnesia
can cause muscle weakness and arrhythmia, but it is a rare condition. Hypomagnesia, defined as low magne-
sium levels in the blood plasma, can be the result of losses in the GI tract, for example, excessive diarrhoea.
Magnesium imbalances can also be a result of alcoholism or secondary to treatment with certain drugs. Hypo-
magnesia is often followed by hypocalcaemia (low calcium ion plasma levels) as well as hypokalaemia and
hyponatraemia [3, 4].

3.3.2 Clinical applications and preparations

Magnesium ion imbalances can manifest in a variety of conditions such as hypo- and hypermagnesaemia.
Magnesium ion preparations are also used as antacids, mostly in combination with aluminium-based salts
(see Section 4.3.5). Additionally, magnesium salts are involved in the treatment of arrhythmia (irregular heart
beat) and eclampsia, a life-threatening hypertensive disorder in pregnant women.

Symptomatic hypomagnesaemia is associated with plasma serum Mg2+ levels of <0.5–1 mmol/kg for a
period of 5 days or more. Mg2+ ions are initially given as intravenous (i.v.) or intramuscular injection; the
latter is fairly painful and consisting of magnesium sulfate (MgSO4). MgSO4 can also be used as emergency
treatment for very serious arrhythmias, a disorder of the heart rate (pulse). In an emergency treatment, it is
usually given intravenously as one single dose or with one repeat (Figure 3.5) [4, 5].

Note that the plasma magnesium concentration should be monitored, and the dose has to be reduced in
patients with renal impairment as Mg2+ is excreted via the kidneys. Magnesium ions can also be given orally
to the patient, for example, in the form of magnesium glycerophosphate tablets.

Magnesium hydroxide [Mg(OH)2] is present in antacids because of its laxative properties and is also the
main ingredient of the ‘milk of magnesia’. The ‘milk of magnesia’ is a suspension of Mg(OH)2 in water,
which has a milk-like appearance because of the low aqueous solubility of Mg(OH)2. It is considered as a
strong electrolyte and a weak base and is given to the patient for indigestion and heartburn. The alkaline
suspension neutralises any excess stomach acid and therefore works as an antacid. It also stimulates intestinal
movement, as the magnesium ions increases the water content in the intestines through its osmotic effect and
as a result softens any faeces present.

Magnesium trisilicate (Mg2Si3O8) can also be used in antacid preparations especially in the treatment of
peptic ulcers. The mode of action includes the increase of the pH of the gastric fluid together with the formation
of a colloidal silica precipitate, which forms a protection for the GI mucosa. Most antacids contain a mixture of

O

OSO

Mg2+
O

Figure 3.5 Structure of MgSO4

Alkaline Earth Metals 55

Mg2+
OO

O Si Si O

OO
Si

OO
Mg2+

Figure 3.6 Chemical structure of magnesium trisilicate (Mg2Si3O8)

aluminium hydroxide [Al(OH)3] and magnesium and/or calcium preparations. Therefore, the mode of action
will be further discussed in the chapter on aluminium-based drugs (see Section 4.3.5) (Figure 3.6).

Unfortunately, orally taken magnesium salts can show interactions with other drugs taken simultaneously.
Magnesium trisilicate reduces the absorption of iron products, certain antibiotics (such as Nitrofurantoin) or
antimalarial drugs (such as Proguanil). Magnesium salt preparations, which form part of antacids, are not
recommended to be taken at the same time as a variety of drugs such as ACE inhibitors, aspirin and peni-
cillamine. In most cases, antacids reduce the absorption of the simultaneously taken drug. Therefore, before
any treatment with antacids, the full medical history of the patient should be taken and possible interactions
assessed [4].

3.4 Calcium: the key to many human functions

Calcium is the most abundant inorganic element in the human body and is an essential key for many physio-
logical processes. Ca2+ has numerous intra and extracellular physiological roles, for example, a universal role

as messenger and mediator for cardiac, skeletal and smooth muscle contractions. Calcium ions are a critical

factor in several life-defining biochemical processes as well as in the endocrine, neural and renal aspects of

blood pressure homeostasis.

Calcium has the symbol Ca and atomic number 20 and is a soft grey alkaline earth metal. Calcium has
four stable isotopes (40Ca and 42Ca–44Ca) and the metal reacts with water with the formation of calcium

hydroxide and hydrogen.

2Ca + 2H2O → 2CaOH + H2 (3.3)

Calcium salts can be found in everyday life. Limestone, cement, lime scale and fossils are only a few
examples where we encounter Ca2+. They also have a wide spectrum of applications spanning from

insecticides to clinical applications. Calcium arsenate [Ca3(AsO4)2] is extremely poisonous and is used in
insecticides. Calcium carbonate (CaCO3) can be found in clinical applications such as antacids, but note
that an excessive intake can be hazardous. Calcium chloride (CaCl2) is used in ice removal and dust control
on dirt roads, as a conditioner for concrete and as an additive in canned tomatoes. Calcium cyclamate

[Ca(C6H11NHSO4)2] is used as a sweetening agent, and calcium gluconate [Ca(C6H11O7)2] is used as a
food additive in vitamin pills. Calcium hypochlorite Ca(OCl)2 can be found in swimming pool disinfectants,
in bleaching agents, in deodorants and in fungicides. Calcium permanganate [Ca(MnO4)2] is used in textile
production, as a water-sterilising agent and in dental procedures. Calcium phosphate [Ca3(PO4)2] finds

56 Essentials of Inorganic Chemistry

applications as a supplement for animal feed, as a fertiliser, in the manufacture of glass and in dental
products. Calcium sulfate (CaSO4⋅2H2O) is the common blackboard chalk.

3.4.1 Biological importance

Calcium ions play important roles in the human body in a variety of neurological and endocrinological
processes. Calcium is known as a cellular messenger and it has a large intra- versus extracellular gradi-
ent (1 : 10 000), which is highly regulated by hormones. This gradient is necessary to maintain the cellular
responsiveness to diverse extracellular stimuli. Calcium ions are also involved in the formation of bones and
teeth, which act also as a reservoir for calcium ions.

A normal adult body contains ∼1000 g of calcium, of which around 99% are extracellular and most of
which is stored in bones and teeth. Bones actually serve as a dynamic store for Ca2+. The remaining 1%
of Ca2+ can be found in the extracellular space, such as plasma, lymph and extracellular water. The intra
and extracellular Ca2+ concentration is extremely important to many physiological functions and is therefore
rigorously controlled (Figure 3.7) [6].

Calcium ions are regulated within the gut, skeleton and kidneys. The Ca2+ homeostasis is normally in equi-
librium, which means that the amount of Ca2+ enters the body is equal to the amount of Ca2+ leaving the body.
Calcium ion levels are regulated by hormones that are not regulated by the Ca2+ level, called noncalciotropic
hormones, for example, sex hormones and growth factors. In contrast, there are hormones that are directly
related to Ca2+, for example, PTH (parathyroid hormone), which are called calciotropic hormones. PTH con-
trols the serum plasma level of Ca2+ by regulating the re-absorption of Ca2+ in the nephron, stimulating the
uptake of Ca2+ from the gut and releasing Ca2+ from the bones which act as a reservoir.

Modified hydroxylapatite, also frequently called hydroxyapatite and better known as bone mineral, makes
up ∼50% of our bones. Hydroxylapatite is a natural form of the mineral calcium apatite, whose formula is

Dietary calcium Extracellular space Bone
Dietary habits, 900 mg
supplements
500 mg
1000 mg

300 mg

125 mg 9 825 mg 500 mg 990 g
Vit D-R 10 000 mg

Vit D

1-α hydroxylase PTH

Small 1,25(OH)2D 25(OH)2D Calcitonin
intestine 175 mg Kidney Estrogens

825 mg

Figure 3.7 Calcium homeostasis [6] (Reproduced with permission from [6]. Copyright © 2005, John Wiley &
Sons, Ltd.)

Alkaline Earth Metals 57

usually denoted as Ca10(PO4)6(OH)2. Modifications of hydroxylapatite can also be found in the teeth, and a
chemically identical substance is often used as filler for replacement of bones, and so on. Nevertheless, despite
similar or identical chemical compositions, the response of the body to these compounds can be quite different.

3.4.2 How does dietary calcium intake influence our lives?

It is believed that an optimal dietary calcium intake can prevent chronic diseases. In the Stone Age, the average
calcium intake was 2000–3000 mg Ca2+/day per adult, whereas now-a-days it has decreased to an average of
600 mg/day [7]. This means that we are living in permanent calcium deficiency, and it is believed that there
are linkages to various chronic diseases, such as bone fragility, high blood pressure and colon cancer [8].

Ca2+ is an essential nutrient, and the required amount varies throughout a person’s life time depending on
the stage of life. There have been three stages of life identified when the human body needs an increased level
of Ca2+. The first one is childhood and adolescence because from birth to the age of ∼18 the bones form
and grow until they reach their maximum strength. Pregnancy and lactation has also been identified as a time
when the human body is in need of an increased level of Ca2+. A full infant accumulates around 30 g of Ca2+
during gestation and another 160–300 mg/day during lactation. Ageing has been identified as the third period
of life in humans when increased calcium intake is required. This has been associated with several changes
to the calcium metabolism in the elderly (Table 3.2).

3.4.3 Calcium deficiency: osteoporosis, hypertension and weight management

Osteoporosis is most commonly associated with calcium deficiency, but an adequate calcium intake should
not only be considered as a therapy for bone loss. It should be seen as an essential strategy for the maintenance

Table 3.2 Optimal daily calcium intake
according to NIH Consensus Conference [6]

Age mg/d

Neonates 400
0–6 mo 600
6–12 mo
800
Children 800 – 1200
1–5 yr
6–10 yr 1200 – 1500

Adolescents 1000
11–24 yr 1500

Male adults 1000
25–65 yr 1200 – 1500
Elderly 1500
1500
Female adults
20–25 yr
Pregnant and nursing
Postmenopausal (>50 yr)
Elderly (>65 yr)

Source: Reproduced with permission from [6]. Copyright ©
2005, John Wiley & Sons, Ltd.

58 Essentials of Inorganic Chemistry

of health in the ageing human. Ninety-nine percent of Ca2+ is found in the bones, as they function as a reser-
voir. Osteoporosis is known to be the major underlying cause on bone fractures in postmenopausal women.
Calcium uptake and plasma concentrations are closely regulated by hormones, as outlined in Section 3.4.1.
Nevertheless, there has been no clear and direct relationship between Ca2+ intake and bone health established
until now. It is believed that a high Ca2+ concentration and vitamin D level is essential in the first three decades
of life in order to establish an optimum bone density level. These also modify the rate of bone loss, which is
associated with ageing.

Studies support the hypothesis that calcium supplementation can reduce blood pressure, being more ben-
eficial to salt-dependent hypertension. The regulation of the cellular calcium metabolism is central to blood
pressure homeostasis. It is believed that the higher the level of cytosolic-free calcium ions, the greater the
smooth muscle vasoconstrictor tone, which in turn has an effect on the sympathetic nervous system activity
and thus on the blood pressure. Nevertheless, studies do not justify the use of calcium supplementation as the
sole treatment for patients with mild hypertension.

It has been hypothesised that there exists a link between dietary calcium and weight management in humans.
It has been proposed that a low-calorie, high-Ca2+ diet helps in supporting the fight against obesity and
increase the energy metabolism. The recommended Ca2+ intake should be around 1200 mg/day as previously
mentioned depending on the age. Available evidence indicates that increasing the calcium intake may sub-
stantially reduce the risk of being overweight, although long-term, large-scale prospective clinical trials need
to be conducted to confirm or better clarify this association.

3.4.4 Renal osteodystrophy

Renal osteodystrophy, also called renal bone disease, is a bone mineralisation deficiency seen in patients with
chronic or end-stage renal failure.

Vitamin D is usually activated in the liver to the pro-hormone calcidiol and then in the kidney to calcitriol,
which is the active form of vitamin D. Both activation steps are based on a hydroxylation reaction. Pro-vitamin
D is hydroxylated in the 25 position in the liver (calcidiol) and then in the kidney at the 1α-position (calcitriol).
Calcitriol helps the body to absorb dietary Ca2+ (Figure 3.8).

In patients with renal failure, the activation to calcitriol is depressed, which results in a decreased con-
centration of Ca2+ in the blood plasma. Furthermore, the plasma phosphate level increases as a result of the
kidney impairment. This, in turn, reduces the amount of free Ca2+ in the blood even more, as the phosphate
complexes the free Ca2+. The pituitary gland senses the low levels of plasma Ca2+ and releases PTH. As
previously outlined, PTH increases the re-absorption of Ca2+ in the nephron and absorption in the gut, and
promotes the release of Ca2+ from the bones. In turn, this leads to a weakening of the bone structure.

Promotes
dietary
uptake of
Ca2+

Vitamin D Liver Calcidiol Kidney Calcitriol

Level too low

Release of
PTH

Figure 3.8 Activation of vitamin D

Alkaline Earth Metals 59

Patients can be treated with phosphate binders in order to avoid excess phosphate absorption from the gut.
Dialysis will also be helpful in removing excess phosphate from the blood. Furthermore, the patient can be
given synthetic calcitriol and potentially calcium supplements.

3.4.5 Kidney stones

Around 20–40% of all kidney stones are associated with elevated Ca2+ level in the urine. For a long time,
it has been suggested that low dietary calcium intake would be the best method to prevent the recurrence of
kidney stones. More recent studies involving patients who suffered from recurring calcium oxalate stones
showed that a low calcium diet did not prevent the formation of kidney stones. It was actually found that
a higher calcium intake of around 1200 mg/day resulted in a significant reduction of the recurrence of kid-
ney stones by around 50%. It is believed that the restriction of calcium leads to an increase in absorption
and excretion of oxalate in the urine and therefore promotes the formation of calcium oxalate stones. Cur-
rently, the conclusion is that kidney stone formation in healthy individuals is not associated with calcium
supplementation [4].

3.4.6 Clinical application

Calcium supplements are usually required only if the dietary Ca2+ intake is insufficient. As previously men-
tioned, the dietary requirements depend on the age and circumstances; for example, an increased need can
be seen in children, in pregnant women and in the elderly where absorption is impaired. In severe acute
hypocalcaemia, a slow i.v. injection of a 10% calcium gluconate has been recommended. It has to be kept
in mind that the plasma Ca2+ level and any changes to the electrocardiogram (ECG) have to be carefully
monitored [5].

A variety of calcium salts are used for clinical application, including calcium carbonate, calcium chloride,
calcium phosphate, calcium lactate, calcium aspartate and calcium gluconate. Calcium carbonate is the most
common and least expensive calcium supplement. It can be difficult to digest and may cause gas in some
people because of the reaction of stomach HCl with the carbonate and the subsequent production of CO2
(Figure 3.9).

Calcium carbonate is recommended to be taken with food, and the absorption rate in the intestine depends
on the pH levels. Taking magnesium salts with it can help prevent constipation. Calcium carbonate consists
of 40% Ca2+, which means that 1000 mg of the salt contains around 400 mg of Ca2+. Often, labels will only
indicate the amount of Ca2+ present in each tablet and not the amount of calcium carbonate (Figure 3.10).

CaCO3 + 2HCI → CO2 + CaCl2 + H2O
Figure 3.9 Chemical equation showing the synthesis of CO2 under acidic stomach conditions

Ca2+ O

OO
Figure 3.10 Chemical structure of calcium carbonate

60 Essentials of Inorganic Chemistry OH
O O

O O O
O Ca2+ Ca2+

Ca2+

O O
O O

OO
OH

Figure 3.11 Chemical structure of calcium citrate

O
O
OH

Ca2+
HO O

O

Figure 3.12 Chemical structure of calcium lactate

Calcium citrate is more easily absorbed (bioavailability is 2.5 times higher than calcium carbonate); it is
easier to digest and less likely to cause constipation and gas than calcium carbonate. Calcium citrate can
be taken without food and is more easily absorbed than calcium carbonate on an empty stomach. It is also
believed that it contributes less to the formation of kidney stones. Calcium citrate consists of around 24%
Ca2+, which means that 1000 mg calcium citrate contains around 240 mg Ca2+. The lower Ca2+ content
together with the higher price makes it a more expensive treatment option compared to calcium carbonate,
but its slightly different application field can justify this (Figure 3.11).

The properties of calcium lactate are similar to those of calcium carbonate [9], but the former is usually more
expensive. Calcium lactate contains effectively less Ca2+ per gram salt than, for example, calcium carbonate.
Calcium lactate consists of only 18% Ca2+, making it a less ‘concentrated’ salt (Figure 3.12) [10].

Calcium gluconate is prescribed as a calcium supplement, but it is also used in the urgent treatment
of hyperkalaemia (K+ plasma levels above 6.5 mmol/l). Hyperkaleamia in the presence of ECG changes
usually requires immediate treatment, and a 10% calcium gluconate solution intravenously administered
is recommended (see Section 2.4.4). Administration of the calcium solution does not lower the plasma
K+ level but protects temporarily against myocardial excitability and therefore temporarily reduces the
toxic effects of hyperkalaemia. Calcium gluconate contains effectively the least Ca2+ per amount of
supplement (only around 9%). That means that in 1000 mg calcium gluconate, only 90 mg is actual Ca2+
(Figure 3.13).

Alkaline Earth Metals 61

OH OH

O
HO

OH OH O Ca2+
OH OH

O
HO

OH OH O
Figure 3.13 Chemical structure of calcium gluconate

3.4.7 Side effects

Several large long-term studies have shown that a daily intake of 1000–2500 mg of calcium salts is safe.
Side effects have been observed only at relatively high doses, being manifested in GI disturbances such as
constipation and bloating and, in extreme cases, arrhythmia [11]. The GI system normally adjusts after a while,
and problems should resolve themselves. Calcium salts are generally better absorbed in an acid environment,
so patients with a low production of stomach acid or elderly patients who are on high doses of antiulcer
medication might experience problems with absorption. It is then recommended to consume the calcium
supplement with a meal [5].

Nevertheless, it is important to note that calcium ions can interfere with the absorption of some drugs, such
as antibiotics. For example, tetracycline and quinolone antibiotics can chelate Ca2+ ions and form complexes
which cannot be absorbed anymore. Therefore, calcium supplements and antibiotics should not be taken
together. Patients are typically advised to take antibiotics 1 h before or 2 h after food [5].

3.5 Barium: rat poison or radio-contrast agent?

The element barium (Ba) has the atomic number 56 and is classified as a heavy metal. Barium metal is highly
reactive and therefore no elemental barium exists in nature. Natural sources of barium are the water-insoluble
minerals barite (barium sulfate) and whiterite (barium carbonate). In order to obtain pure barium compounds,
the mineral barite is reacted with carbon, and barium sulfide is formed. Barium sulfide is, in contrast to
barium sulfate, water soluble. Subsequently, the pure barium sulfide is treated with sulfuric acid and pure
barium sulfate can be obtained.

BaSO4 + 4C → BaS + 4CO

BaS + H2SO4 → BaSO4 + H2S

Barium salts can be highly toxic even at low concentrations. Barium carbonate is highly toxic and can
be used as rat poison as it readily dissolves in the stomach acid. Barium sulfate is the least toxic barium
compound mainly because of its insolubility. Barium sulfate is used in a variety of applications ranging from
white paint to X-ray contrast agent.

62 Essentials of Inorganic Chemistry

O
Ba2+

S
OO

O

Figure 3.14 Chemical structure of barium sulfate

The clinical use of barium sulfate suspension is well known under the term barium meal. Patients are given
a suspension of barium sulfate to swallow. Using X-ray imaging, the whole oesophagus, the stomach and the
intestines can be visualised. Barium sulfate lines the tissue whilst travelling through the digestive tract. The
heavy barium ions absorb X-rays readily and therefore these structures become visible in an X-ray screening.
Barium sulfate is a well-used and tolerated oral radio-contrast agent. It is also used as radio-contrast agent in
enemas (Figure 3.14) [3, 4].

Alkaline Earth Metals 63

3.6 Exercises

3.6.1 Calcium supplementation
Calcium supplementation is recommended as a dietary supplement especially for menopausal, preg-
nant or nursing women. There are a variety of calcium salts on the market that can be used for oral
administration. For the examples given below, determine the chemical formula, the molecular weight
and the Ca2+ content (expressed in gram/gram (g/g) and percentage weight/weight (%w/w)).

(a) Calcium carbonate
(b) Calcium lactate
(c) Calcium chloride
(d) Calcium citrate
(e) Calcium gluconate

3.6.2 Complete the redox equation (including the half-equations) and indicate the standard reduction
3.6.3 potential assuming that both reaction partners are present in the same concentration.

(a) Mg + Cl2 → ??
(b) Ba + Br2 → ??

Complete the following redox equation (including the half-equations).

Mg + Cl2 →??

Indicate the standard reduction potential assuming a concentration of [Mg2+] = 0.7 mol/l and
[Cl−] = 0.8 mol/l.

3.6.4 Milk of magnesia
Milk of magnesia is typically an 8.7% w/v aqueous suspension of magnesium hydroxide.

(a) What are the chemical formula and the molecular weight of magnesium hydroxide?
(b) What is the concentration of magnesium hydroxide in gram per litre?
(c) How many moles of such magnesium salts are present in a 100-ml suspension?



Alkaline Earth Metals 65

3.7 Case studies

3.7.1 Magnesium hydroxide suspension

Magnesium hydroxide mixture is an aqueous oral suspension containing hydrated magnesium oxide. It is
indicated for use in constipation in adults and children. Typical analysis methods used for quality purposes
are based on titration reactions. A certain volume of the suspension containing hydrated magnesium oxide
[Mg(OH)2] is typically reacted with a known amount of sulfuric acid (H2SO4). The excess acid is then titrated
with sodium hydroxide (NaOH) and methyl orange as an indicator [12].
(a) Research the type of titration described. Describe the chemical structure and mode of action of the

indicator.
(b) Formulate the relevant chemical equations.
(c) For the analysis, 10 g of the suspension was reacted with 50 ml of 0.5 M H2SO4. The excess H2SO4 was

titrated with 1 M NaOH using methyl orange as indicator. For each titration, the following volume of
NaOH has been used:

11.0 ml 11.2 ml 10.9 ml

Calculate the amount of Mg(OH)2 present in your sample. Express your answer in grams and moles.

3.7.2 Calcium carbonate tablets

Your pharmaceutical analysis company has been contacted by an important client and asked to analyse a
batch of injections containing calcium carbonate (CaCO3). The description of your brief states that you are
supposed to analyse the active pharmaceutical ingredient (API) in these tablets following standard quality
assurance guidelines.

Typical analysis methods used for quality purposes are based on titration reactions. A certain amount of
the tablet powder is dissolved in water and hydrochloride acid (HCl). A known amount of disodium edetate
is added. After adjustment of the pH, the excess disodium edetate is titrated with zinc chloride (ZnCl) using
morbant black II solution as indicator [12].
(a) Research the type of titration described. Describe the chemical structure and mode of action of the

indicator. You may want to familiarise yourself with chelation (see Section 11.2).
(b) Formulate the relevant chemical equations.
(c) The package states that each tablet contains 1.5 g of CaCO3. For the experiment, 20 tablets are weighed

(total weight 42.6 g) and powdered. An amount of powder containing 50 mg of Ca2+ is dissolved in
water and HCl and reacted with 50 ml of 0.05 M disodium edetate. After adjusting the pH to 10.9, the
excess disodium edetate is titrated with 0.05 M ZnCl2 solution. For each titration, the following volume
of ZnCl2 has been used:

25.0 ml 24.8 ml 25.3 ml

Calculate the amount of CaCO3 present in your sample. Express your answer in grams and moles.
(d) Critically discuss your result in context with the stated value for the API.
(e) Research the typically accepted error margins.

66 Essentials of Inorganic Chemistry

References

1. J. D. Lee, Concise inorganic chemistry, 5th ed., Chapman & Hall, London, 1996.
2. B. P. Barna, D. A. Culver, B. Yen-Lieberman, R. A. Dweik, M. J. Thomassen, Clin. Diagn. Lab. Immunol. 2003, 10,

990 – 994.
3. G. J. Tortora, B. Derrickson, Principles of anatomy and physiology, 12th ed., international student/Gerard J. Tortora,

Bryan Derrickson. ed., Wiley [Chichester: John Wiley, distributor], Hoboken, N.J., 2009.
4. G. A. McKay, M. R. Walters, J. L. Reid, Lecture notes. Clinical pharmacology and therapeutics, 8th ed.,

Wiley-Blackwell, Chichester, 2010.
5. British national formulary, British Medical Association and Pharmaceutical Society of Great Britain, London.
6. E. R. Tiekink, M. Gielen, Metallotherapeutic drugs and metal-based diagnostic agents: the use of metals in medicine,

Wiley, Chichester, 2005.
7. S. B. Eaton, D. A. Nelson, Am. J. Clin. Nutr. 1991, 54, S281–S287.
8. M. J. Bargerlux, R. P. Heaney, J. Nutr. 1994, 124, S1406–S1411.
9. B. R. Martin, C. M. Weaver, R. P. Heaney, P. T. Packard, D. L. Smith, J. Agric. Food Chem. 2002, 50, 3874–3876.
10. D. A. Straub, Nutr. Clin. Pract. 2007, 22, 286–296.
11. (a) M. A. Oconnell, J. S. Lindberg, T. P. Peller, H. M. Cushner, J. B. Copley, Clin. Pharm. 1989, 8, 425–427; (b) W.

F. Caspary, Eur. J. Gastroen. Hepat. 1996, 8, 545–547.
12. British pharmacopoeia, Published for the General Medical Council by Constable & Co, London.

Further Reading

1. E. Alessio, Bioinorganic medicinal chemistry, Wiley-VCH, Weinheim, 2011.
2. W. Kaim, B. Schwederski, Bioinorganic chemistry: inorganic elements in the chemistry of life: an introduction and

guide, Wiley, Chichester, 1994.
3. H.-B. Kraatz, N. Metzler-Nolte, Concepts and models in bioinorganic chemistry, Wiley-VCH [Chichester: John

Wiley, distributor], Weinheim, 2006.
4. R. M. Roat-Malone, Bioinorganic chemistry: a short course, Wiley, Hoboken, N.J. [Great Britain], 2002.

4

The Boron Group – Group 13

Group 13 (13th vertical column of the periodic table) is called the boron group and it consists of boron (B),
aluminium (Al), gallium (Ga), indium (In) and thallium (Tl) (Figure 4.1).

All elements within group 13 show a wide variety of properties. It is important to note that boron is a metal-
loid (semi-metal) whereas aluminium is a metal but shows many chemical similarities to boron. Aluminium,
gallium, indium and thallium are considered to be metals of the ‘poor metals’ group.

Metalloids are elements that display some properties characteristic for metals and some characteristic
for nonmetals.

In this chapter, the general chemistry of group 13 elements is discussed as well as some clinical applications
for boron and aluminium. Further clinical applications for boron as well as applications for thallium can be
found in the chapter on radiochemistry (Chapter 10).

4.1 General chemistry of group 13 elements

Group 13 elements are characterised by having three electrons in their valence shell. Therefore, all elements
form the stable cation M3+. Most elements (with the exception of B) form additionally the singly positively
charged ion M+, which is indeed the more stable oxidation state for Tl.

Boron and aluminium occur only with oxidation number +3 in their compounds, and with a few exceptions
their compounds are best described as ionic. The electronic configuration shows three electrons outside a noble
gas configuration, two in an s shell and one in a p shell. The outermost p electron is easy to remove as it is
furthest from the nucleus and well shielded from the effective nuclear charge. The next two s electrons are also
relatively easy to remove. Removal of any further electrons disturbs a filled quantum shell and is therefore
difficult. This is reflected in the ionisation energies (Table 4.1).

The main sources of B are the two minerals borax (Na2[B4O5(OH)4]⋅8H2O) and kernite (Na2[B4O5(OH)4]),
which are generally used as components in many detergents or cosmetics. Al occurs widely on earth, and it

Essentials of Inorganic Chemistry: For Students of Pharmacy, Pharmaceutical Sciences and Medicinal Chemistry,
First Edition. Katja A. Strohfeldt.
© 2015 John Wiley & Sons, Ltd. Published 2015 by John Wiley & Sons, Ltd.
Companion website: www.wiley.com/go/strohfeldt/essentials

68 Essentials of Inorganic Chemistry

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La- Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Lu

Fr Ra Ac- Rf Db Sg Bh Hs Mt Ds Rg Uub
Lr

Figure 4.1 The periodic table of elements, group 13 elements are highlighted

Table 4.1 Ionisation energy (kJ/mol) for
group 13 elements [1]

First Second Third

B 801 2427 3659
Al 577 1816 2744
Ga 579 1979 2962
In 558 1820 2704
Tl 589 1971 2877

Source: Reproduced with permission from [1]. Copyright ©
1996, John Wiley & Sons, Ltd.

is the most abundant metal and the third most abundant element in the earth’s crust. Aluminosilicates, such
as clays, micas, feldspar, together with bauxite, are the main sources of Al. Ga, In and Tl occur in traces as
their sulfides.

4.1.1 Extraction

Boron (B) can be extracted from borax by converting the latter to boric acid (Equation 4.1) and subsequently
to the corresponding oxide (Equation 4.2). Boron of low quality can then be obtained by the reduction of
boron oxide with Mg, followed by several steps of washing with bases and acids.

Na2[B4O5(OH)4] ⋅ 8H2O + H2SO4 → 4B(OH)3 + Na2SO4 + 5H2O (4.1)
2B(OH)3 → B2O3 + 3H2O (4.2)

Al is extracted from ores such as bauxite or cryolite in the so-called Bayer process. Bauxite contains mainly

a mixture of aluminium oxides with Fe2O3, SiO2 and TiO2 as impurities. In the Bayer process, hot aqueous
NaOH is added to the crude ore under pressure and aluminium hydroxide will go into solution. This will result
in the separation of Fe2O3. The solution is cooled down and seeded with Al2O3⋅3H2O in order to precipitate

The Boron Group – Group 13 69

NaOH AI(OH)3

Bauxite Na[AI(OH)4] Crystallisation

Electrolysis

Pure AI AI2O3

Figure 4.2 Bayer process

Al(OH)3. Pure Al can be produced by electrolysis of molten Al2O3 (melting point 2345 K), with Al being
obtained at the cathode (Figure 4.2).

The main source of Ga is bauxite, but it can also be obtained from the residues from the Zn processing

industry. It can be found in the zinc sulfide ore sphalerite. Tl can be obtained as by-product of the processing
of Cu and Zn ores. The demand for In and Tl is rather low.

4.1.2 Chemical properties

4.1.2.1 Reactivity

B is chemically unreactive except at high temperatures. Al is a highly reactive metal, which is readily oxi-

dised in air to Al2O3. This oxide coating is resistant to acids but is moderately soluble in alkalis. Al itself
dissolves in diluted mineral acids (Equation 4.3) and can react with strong alkalis, the product being the
tetrahydroxoaluminate ion [Al(OH)4−] and H2 (Equation 4.4).

Al + 3H2SO4(aq) → Al(SO4)3 + 3H2 (4.3)
2Al + 2MOH + 6H2O → 2M[Al(OH)4] + 3H2 (4.4)

Aluminium can be used to reduce metal oxides, the most famous example being the thermit process. Al
reacts violently with iron(Ill) oxide to produce iron in this highly exothermic process, where Fe is obtained
in its liquid form (Equation 4.5).

2Al(s) + Fe2O3(s) → 2Fe(l) + Al2O3(s) (4.5)

Ga, In and Tl dissolve in most acids, and as a result the salts of Ga(III), In(III) and Tl(I) are obtained,
whereas only Ga reacts with aqueous alkali with the production of H2.

4.1.2.2 Oxides/hydroxides: amphoteric compounds

Boron oxide (B2O3) is an acidic oxide and an insoluble white solid with a very high boiling point (over 2000 K)
as a result of its extended covalently bonded network structure. Aluminium oxide (Al2O3, Equations 4.6 and
4.7) as well as aluminium hydroxide (Al(OH)3, Equations 4.8 and 4.9) are amphoteric compounds.

Amphoteric compounds are substances that can react either as an acid or as a base.

70 Essentials of Inorganic Chemistry

Al2O3 + 3H2O + 2(OH−) → 2[Al(OH)4]− (4.6)

Al2O3 + 3H2O + 6[H3O]+ → 2[Al(H2O)6]3+ (4.7)

Al(OH)3 can neutralise a base and therefore act as an acid (Equation 4.8); it can also neutralise an acid and
act as a base (Equation 4.9).

Al(OH)3 + NaOH → NaAl(OH)4 (4.8)
Al(OH)3 + 3HCl → AlCl3 + 3H2O (4.9)

4.1.2.3 Halides

The most important halide of boron is the colourless gas boron trifluoride (BF3). Aluminium chloride (AlCl3)
is a volatile solid which sublimes at 458 K. The vapour formed on sublimation consists of an equilibrium
mixture of monomers (AlCl3) and dimers (Al2Cl6). It is used to prepare the powerful and versatile reducing
agent lithium tetrahydridoaluminate (LiAlH4).

Both boron trichloride (BCl3) and aluminium trichloride (AlCl3) act as Lewis acids to a wide range of
electron-pair donors, and this has led to their widespread use as catalysts. In the important Friedel–Crafts
acylation, AlCl3 is used as a strong Lewis acid catalyst in order to achieve the acylation of an aromatic ring.

A Lewis acid is defined as a compound that can accept electrons pairs with the formation of a coordinate

covalent bond. Any type of electrophile can be a Lewis acid. In contrast, Brønsted–Lowry acids are com-
pounds that transfer a hydrogen ion (H+) and they are the more commonly known type of acids. Analogous
definitions apply for a Lewis base (electron donator) and a Brønsted–Lowry base (H+ acceptor).

The following sections will describe the clinical application of boron, aluminium and gallium. It is impor-
tant to note that more information on the clinical use of gallium and thallium can be found in Section 10.4,
where radiopharmaceuticals are discussed.

4.2 Boron

4.2.1 Introduction

Boron has the atomic number 5 and the symbol B, and is a so-called metalloid (see Chapter 4). Boron
compounds have been known for many centuries and especially used in the production of glass. Boric acid
[B(OH)3] is used in the large-scale production of glass. Borosilicate glasses (Pyrex® glass), which are pro-
duced by a fusion of B2O3 and silicate, are extremely heat resistant and often used in laboratories.

At the beginning of the nineteenth century, it was recognised that boron is an essential micronutrient for
plants. A deficiency of boron can lead to deformation in the vegetable growth such as hollow stems and hearts.
Furthermore, the plant growth is reduced and fertility can be affected. In general, boron deficiency leads to
qualitative and quantitative reduction in the production of the crop. Boron is typically available to plants as
boric acid [B(OH)3] or borate [B(OH)4]−. The exact role of boron in plants is not understood, but there is
evidence that it is involved in pectin cross-linking in primary cell walls, which is essential for normal growth
and development of higher plants [2].

Borax (Na2[B4O5(OH)4]⋅8H2O) can be applied as a fertiliser and, together with kernite (Na2[B4O5(OH)4]⋅
2H2O), forms the two most commercially available borates. Borates find a wide range of practical applications

The Boron Group – Group 13 71

H
O

HBH
OO

Figure 4.3 Chemical structure of boric acid

such as in detergents, cosmetics, antifungal mixtures as well as components in fibreglass and others. The
toxicity of borates in mammals is relatively low, but it exhibits a significantly higher risk to arthropods and
can be used as an insecticide.

Boron-based compounds are used in a wide range of clinical applications including their use as antifun-
gal and antimicrobial agent, as proteasome inhibitors and as a noninvasive treatment option for malignant
tumours. The latter application will be discussed in the chapter on radiopharmaceuticals (Chapter 10).

4.2.2 Pharmaceutical applications of boric acid

Boric acid is a long-standing traditional remedy with mainly antifungal and antimicrobial effects. For medic-
inal uses, it has become known as sal sedativum, which was discovered by Homberg, the Dutch natural
philosopher, in 1702 [3]. Diluted solutions were and sometimes still are used as antiseptics for the treatment
of athletes’ foot and bacterial thrush, and in much diluted solutions as eyewash (Figure 4.3) [4].

Boric acid can be prepared by reacting borax with a mineral acid:

Na2B4O7 ⋅ 10H2O + 2HCl → 4B(OH)3[or H3BO3] + 2NaCl + 5H2O
In general, there are many other health claims around the clinical use of boric acid and boron-containing
compounds, but many of those have no supporting clinical evidence.

4.2.3 Bortezomib

Bortezomib belong to the class of drugs called proteasome inhibitors and is licensed in the United States and
the United Kingdom for the treatment of multiple myeloma. The drug has been licensed for patients in whom
the myeloma has progressed despite prior treatment or where a bone marrow transplant is not possible or was
not successful. It is marketed under the name Velcade® or Cytomib®. Velcade is administered via injection
and is sold as powder for reconstitution (Figure 4.4) [5].

Bortezomib was the first drug approved in the new drug class of proteasome inhibitors and boron seems to
be its active element. For the mode of action, it is believed that the boron atom binds with high affinity and
specificity to the catalytic site of 26S proteasome and inhibits its action. Therapy with Bortezomib can lead
to a variety of adverse reactions, including peripheral neuropathy, myelosuppression, renal impairment and
gastrointestinal (GI) disturbances together with changes in taste. Nevertheless, the side effects are in most
cases less severe than with alternative treatment options such as bone marrow transplantation [5].

4.3 Aluminium

4.3.1 Introduction

The element aluminium has the atomic number 13 and chemical symbol Al. Aluminium forms a diag-
onal relationship with beryllium. The name ‘aluminium’ derives from the salt alum (potassium alum,

72 Essentials of Inorganic Chemistry

O OH

H
N NB

N OH
H

O
N

Figure 4.4 Chemical structure of bortzemib

KAl(SO4)2⋅12H2O), which was used for medicinal purposes in Roman times. Initially, it was very difficult
to prepare pure aluminium and therefore it was regarded as a very precious substance. In the mid-1800s,
aluminium cutlery was used for elegant dinners, whereas it is nowadays used as lightweight camping cutlery.
In 1886, the manufacture of aluminium by electrolysis of bauxite started, and the price for pure aluminium
dropped significantly. Aluminium is a soft, durable and lightweight metal, which makes it attractive to many
applications. Nowadays, aluminium is mainly used for the construction of cars and aircrafts and can be found
in packaging and construction materials.

4.3.2 Biological importance
The human body contains around 35 mg of Al3+, of which ∼50% is found in the lungs and ∼50% in the
skeleton. There is no known biological role for Al3+ and, indeed, the human body has developed very effective
barriers to exclude it. Only a minimal fraction of Al3+ is taken up from the diet in the gut, and the kidneys
fairly quickly excrete most of it. The bones can act as a sink for Al3+ if the blood concentration is high and
release it slowly over a long period. The brain is vulnerable to Al3+ and usually the blood–brain barrier
prevents Al3+ entering the brain. Al3+ can sometimes act as a competitive inhibitor of essential elements
such as Mg2+, Ca2+ and Fe2+/3+ because of their similar ionic radii and charges. It is important to note that at
physiological pH, Al3+ forms a barely soluble precipitate Al(OH)3, which can be dissolved by changing the
pH (see Equations 4.8 and 4.9) [6].

A normal adult diet contains typically between 2.5 mg/day and up to 13 mg/day Al3+, but patients
on aluminium-containing medication can be exposed to more than 1000 mg/day. Typically, ∼0.001% is
absorbed in the digestive tract, but it can be around 0.1–1.0% when it is in the form of aluminium citrate
(Figure 4.5) [6b].

Al3+ can accumulate in the human body if natural limits are crossed, for example, intravenous admin-
istration or patients on dialysis, or when the kidneys are impaired and therefore not able to excrete Al3+
sufficiently. Under normal circumstances, Al3+ would not accumulate in the human body. Nevertheless, in
1972, Alfrey et al. described the new syndrome of progressive dialysis encephalopathy, the so-called dialysis
dementia, which was seen in patients being treated with haemodialysis for 15 months or more. The symptoms
include speech disorders, problems with the bone mineralisation and general signs of dementia. Investigations
showed that brain scans were normal and that there was no connection to the Alzheimer’s disease, as neither
neurofibrillary tangles nor senile plaques were found. Increased serum and bone concentrations of Al3+ were

The Boron Group – Group 13 73

AI

OO O
OH O

OO
Figure 4.5 Chemical structure of aluminium citrate

found in patients who were on haemodialysis, and the connection was made to the toxicity of the Al3+ present
in the dialysate solution. Nowadays, the use of modern Al3+-free dialysate solutions or new techniques (e.g.

reverse osmosis) prevents ‘dialysis dementia’ [6a].

4.3.3 Al3+ and its use in water purification

Al3+ is used in the purification of water. Lime (CaO) and aluminium sulfate Al2(SO4)3 are added to waste
water in order to accelerate the settling or sedimentation of suspended matter [7]. The addition of lime

increases the pH of the water slightly (Equation 4.10). The water becomes more basic, which promotes the
precipitation of Al3+ as Al(OH)3 (Equation 4.11).

CaO(s) + H2O(l) → Ca2+(aq) + 2OH−(aq) (4.10)
Al3+(aq) + 3OH−(aq) → Al(OH)3(s) (4.11)

Al(OH)3 precipitates as a gelatinous precipitate which slowly settles. During this process, it incorporates
suspended soil, colloidal material and most bacteria. The water is filtered before leaving the treatment plant
in order to remove the flocculate and the vast majority of the Al3+. WHO guidelines allow a maximum con-
centration for Al3+ of 0.2 mg/l [8].

4.3.4 Aluminium-based adjuvants

An adjuvant is an agent or a mixture of agents that possesses the ability to bind to a specific antigen. Adju-
vants are added to vaccines in order to increase the antibody responses to the vaccination and/or to stabilise
the preparation. Adjuvants can absorb many antigenic molecules over a wide surface area, thus enhancing the
interaction of immune cells with the presenting antigens and leading to an increase of the immune response
stimulation. Some adjuvants (including aluminium-based ones) can function as a slow-release delivery sys-
tem. They trap the antigen in a depot created by the adjuvant at the injection site. From there, the antigen is
slowly released, which causes a steady stimulation of the immune system.

Aluminium-based adjuvants have a long-standing tradition and have been used for more than 50 years.
They are the most widely used adjuvants in human and veterinary vaccines and regarded as safe if applied
correctly. Al3+ salts are the only kind of adjuvant licensed by the FDA. They are also the only kind of adju-
vants used in anthrax vaccines for humans in the United States. Anthrax vaccine contains Al(OH)3, as do the
FDA-licensed diphtheria, haemophilus influenzae type B, hepatitis A, hepatitis B, Lyme disease, pertussis
and tetanus vaccines. In many countries, vaccines for children contain aluminium-based adjuvants [9].

The adjuvant effect of potassium alum (KAl(SO4)2⋅12H2O) was first discovered in 1926. Researchers
examined diphtheria toxoids precipitated with alum and were able to show that an injection of this alum
precipitate led to a significant increase in immune response. Leading on from this research, alum has found

74 Essentials of Inorganic Chemistry

widespread use as an adjuvant. Vaccines containing alum as adjuvant are referred to as alum-precipitated
vaccines. Unfortunately, it has been shown that alum precipitations can be highly heterogeneous. The homo-
geneity of the preparation depends on the anions and the conditions present at the point of precipitation [9].

Subsequent research showed that aluminium hydroxide (Al(OH)3) hydrogels can be pre-formed in a stan-
dardised manner and be used to absorb protein antigens to form a homologous preparation. Following on
from this research, researchers have shown that it is possible to co-precipitate aluminium phosphate (AlPO4)
and the diphtheria toxoid in order to form active vaccines. These vaccines are called aluminium-absorbed
vaccines and, in contrast to alum-precipitated vaccines, the antigens are distributed homogeneously. Nowa-
days, aluminium-absorbed vaccines have taken over from alum-precipitated ones. Nevertheless, there is a lot
of ambiguity found in the literature, where both terms are interchangeably used [9].

In summary, immunisation vaccines containing adjuvants are more effective than those without them. Typ-
ical adjuvants are alum [KAl(SO4)2⋅12H2O], Al(OH)3, AlPO4, Al2O3, but oxides of other metals, such as
ZrO2, SiO2 and Fe2O3, are also under investigation.

The formation of the aluminium hydrogels is generally achieved by reacting Al3+ ions (from compound
such as AlCl3) under alkaline aqueous conditions. Conditions are strongly regulated, as even smallest changes
to parameters such as temperature, concentration and others can influence the quality of the hydrogel. Alu-
minium phosphate gels are typically produced by reacting Al3+ salts in the presence of phosphate ions under
alkaline conditions [9].

The mode of action is highly complex and still not fully understood. Initial theories included the physical
absorption of the antigen, which is still considered as an important feature, and the gradual release of antigen
from the injection side with the adjuvant working as an agglomeration. The latter theory was disproved
quickly. Research has shown that antigens need to be adsorbed to the adjuvant before the immunisation
reaction. It is believed that the adjuvant will then present the antigen to the immunocomponent of the targeted
cell [9].

4.3.5 Antacids

The function of antacids is to neutralise excess stomach acid. They also exhibit cytoprotective effects towards

attacks against the gastric mucosa. They are additionally known to heal gastric and duodenal ulcerations;

nevertheless, the mechanism is still uncertain.

Antacids have been in use for the past 2000 years, and the initial formulations were based on CaCO3 (coral
and limestone). Nowadays, the antacid/anti-gas market is a significant income stream for the pharmaceutical

industry and the demand for antacids is expected to grow. The number of people suffering from heartburn

increases with an ageing population, more stressful lifestyles and changing eating habits such as eating out

more often.

Aluminium hydroxide (Al(OH)3) has several medical applications. It is used as an antacid for treating
heartburn as well as acid indigestion (reflux oesophagitis). It is also known to have healing properties of

peptic ulcers. In patients suffering from kidney failure, who show elevated serum phosphate levels (hyper-

phosphataemia), Al(OH)3 is used as a phosphate binder (see Section 4.3.7).
Al(OH)3 is an amphoteric compound (see Section 4.1.2.2), which means it can react as a base or as an

acid. In its application as an anti-acid, Al(OH)3 reacts with any excess stomach acid (mainly HCl) with the
formation of AlCl3 and water (Equation 4.12).

Al(OH)3 + 3HCl → AlCl3 + 3H2O (4.12)

Al(OH)3 is known to cause constipation, so formulations of anti-acids often include a combination with
Mg2+ antacids. Usually, oral antifoaming agents, such as simethicone, are added in order to reduce bloating

The Boron Group – Group 13 75

Table 4.2 Typical formulation of an antacid/antigas
mixture (maximum strength Maalox®, Max®, Norvatis)

Active ingredient Quantity (mg) Purpose

Al(OH)3 400 Antacid
Mg(OH)2 400 Antacid
Simethicone Anti-gas
40

HO O NH2
AI

OH O
Figure 4.6 Chemical structure of dihydroxy aluminium glycinate

and discomfort/pain. Simethicone is a mixture of poly(dimethyl siloxane) and silica gel, which decreases the

surface tension of gas bubbles (Table 4.2).

Ancient anti-acid formulations contained sodium bicarbonate (baking soda, NaHCO3), which resulted in
a rapid reaction with the gastric acid. The result was an increase in the gastric pH and the production of CO2
gas as a by-product (Equation 4.13). Large doses of NaHCO3 can cause alkaline urine and this can result in
kidney problems. Acid neutralisation using Al(OH)3 does not produce CO2 and therefore these side effects
can be avoided.

NaHCO3 + HCl → NaCl + H2O + CO2 (4.13)

Aluminium glycinate [Al(NH2CH2COO)(OH)2] (Figure 4.6) is also used in anti-acid formula-
tions. For example, Gastralgine® contains, amongst other ingredients, dihydroxy aluminium glycinate

[Al(NH2CH2COO)(OH)2], Al(OH)3, magnesium trisilicate and simethicone. It is known to have additionally
protective effects from ulcers.

4.3.6 Aluminium-based therapeutics – alginate raft formulations

Heartburn is the primary symptom of the so-called gastro-oesophageal reflux disease (GERD), which is
caused by the oesophageal influx of gastric HCl from the stomach. There are also close links to oesophageal
cancer, which has a very low survival rate. Relief can be achieved with the use of alginate raft formulations,
which typically contain alginic acid, NaHCO3, magnesium trisilicate and Al(OH)3. Alginates are natural
polysaccharide polymers which are isolated from brown seaweeds.

In the acidic stomach, alginate salts and alginic acids precipitate to form a low-density viscous gel. When
the mixture comes into contact with gastric HCl, the gel matrix formation occurs. HCO3−, which is trapped
in the gel, leads to the formation of CO2 gas (Equation 4.13). The gas bubbles trapped in the gel convert it
to foam and provide buoyancy, allowing the gel to float on the surface of stomach contents (like a raft on
water). Al(OH)3 provides an additional capacity to neutralise any excess stomach acid (Equation 4.12). The
raft physically acts as a barrier to gastric reflux and moves into the oesophagus during reflux. It acts as mobile
neutralising sealant in the oesophageal space when the gastric pressure is high. Once the pressure reduces,
the raft drops back into the stomach and can be digested (Figure 4.7).

76 Essentials of Inorganic Chemistry

Oesophagus

Upwards movement Raft
of raft through CO2 CO2
production
CO2 pocket
Acid

Stomach

Figure 4.7 Illustration of the stomach, showing the acid pocket and the alginate raft floating on top of it protecting
the oesophagus

4.3.7 Phosphate binders

Hyperphosphataemia, that is, increased levels of serum phosphate, is a disorder commonly seen in patients
with end-stage renal (kidney) disease where the kidneys are not able to excrete excess phosphate as a result
of a low renal clearance rate. This disorder is often seen in patients who are on dialysis treatment. Persistent
hyperphosphataemia results in renal osteodystrophy, that is, the weakening of bones due to disturbances in
the calcium and phosphate metabolism.

Generally Al3+-containing drugs are given in order to promote the binding of phosphate in the gut. Antacids
containing Al(OH)3 can be used as phosphate binders. When Al(OH)3 enters the acidic stomach (pH ∼ 1),
Al3+ ions are formed. Some Al3+ ions will be absorbed in the stomach, but the majority is passed to the
distal intestines, where the pH is significantly increased (pH 6–8.5). In this high pH range, Al3+ freshly
precipitates as a colloidal, amorphous Al(OH)3. Its large surface area adsorbs phosphate ions (usually in
form of HPO42−) and passes them through the remaining intestine without decomposition, as the pH is too
high. The Al3+-phosphate complex (AlPO4) is then excreted via the faeces.

Aluminosilicates can also be used as a phosphate binder and is, for example, the active ingredient in
Malinal®. In contrast to Al(OH)3, which acts as an efficient PO43− binder directly, aluminosilicates need
prior exposure to a acid in order to produce free Al3+. Once the free Al3+ is formed, it follows the same
mode of action.

Initially, aluminium-based phosphate binders were also used in dialysis exchange fluids, especially in
patients being treated with haemodialysis. Nevertheless, as a result of the exposure to high concentrations of
Al3+ salts, relatively high concentrations were found in patients. A significant number of patients developed
dementia symptoms after 15 or more months of treatment, which was linked to the high Al3+ concentrations
in the body including the brain (see Section 4.3.2) [6a, 10].

4.3.8 Antiperspirant

Aluminium trichloride (AlCl3) was the first compound that was used as an antiperspirant. The mechanism
of action is still under investigation, but it appears to act by forming a plug of Al(OH)3 within the sweat

The Boron Group – Group 13 77

duct. AlCl3 is a very strong antiperspirant and only advised by doctors if normal antiperspirants do not work.
Leading brands of antiperspirants contain usually a ∼20% aluminium hexahydrate solution in an alcoholic
base. It is thought to work by blocking the openings of the sweat ducts. It tends to work best in the armpits.
However, it may also work for sweating of the palms and soles. It can also be applied to the face, taking care
to avoid the eyes.

4.3.9 Potential aluminium toxicity

The excessive use of aluminium preparations negatively influences human health. Excessive intake of Al3+
has been found to accumulate in sensitive loci and can lead to pathological aberrations and result in dialysis
dementia or similar symptoms. It is important to note that Al3+ is a major component in over-the-counter
drugs such as antacids. Special attention has to be given by the dispensing pharmacist, and the patient has
to be made aware of the consequences of overdoses of Al3+-containing products. Al3+ is known to have
embryonic and foetal toxic effects in humans and animals, causing osteomalacia, which is the softening of
the bones due to defective bone mineralisation [5].

Albumin and transferrin bind around 95% of serum aluminium, which is then cleared mainly via the kidneys
(a small amount can be found in the faeces). In healthy humans, only 0.3% of orally administered aluminium is
absorbed, whereas it has the potential to accumulate when the GI tract is bypassed, for example, in intravenous
infusions [10].

4.4 Gallium

4.4.1 Introduction

Gallium has atomic number 31 in the periodic table of elements. It has a silvery-white colour with a melting
point of only 29 ∘C, which means that it melts when held in the hand. It has no known physiological role in
the human body, but it can interact with cellular processes and proteins that are normally involved in iron
metabolism.

Gallium tartrate has a long research history. Researchers showed in the 1930s that it could be used to
treat syphilis in rabbits with no significant toxicity [6a]. In subsequent studies, it has been shown that gal-
lium ions predominantly accumulate in the bone and therefore would be a good candidate for radiotherapy
of bone cancer. Unfortunately, the radioactive isotope 72Ga has only a half-life of around 14 h, which is
not long enough for effective radiotherapy. Nevertheless, current clinical developments involve the use of
radioactive gallium isotopes as tumour imaging reagents (see Chapter 10), gallium nitrate in metabolic bone
disease, hypercalcaemia and as anticancer drug, as well as up-to-date research in the area of chemotherapeutic
applications.

4.4.2 Chemistry

Gallium exists as the trivalent cation Ga3+, and in aqueous solution it presents as a hydrated complex. Depend-
ing on the pH, a variety of hydroxyl species are formed, some of which are insoluble, such as Ga(OH)3. At
physiological pH, nearly no free Ga3+ is present and the hydroxyl species Ga(OH4)− (gallate, the domi-
nant species) and Ga(OH)3 are formed. Gallium hydroxide species are amphoteric, analogous to aluminium
hydroxide compounds.

It is important to note that the stability of solutions containing gallium chloride or gallium nitrate for oral
administration is affected by the pH. They might not be stable over extended periods and gallium hydroxide
precipitates.

78 Essentials of Inorganic Chemistry

4.4.3 Pharmacology of gallium-based drugs

Ga3+ has an ionic radius and binding properties similar to those of Fe3+ (ferric iron). Unlike Fe3+, it cannot
be reduced to its divalent state, which means that it follows a completely different redox chemistry compared
to iron. The oxidation and reduction of iron is important in many biological processes, which therefore cannot
be mimicked by gallium. One example includes the uptake of Fe2+ by the haeme group (see Chapter 8). As
Ga3+ is not readily reduced to it +II state, it cannot bind to the haeme group.

Transferrin is an important transport protein that controls the level of free Fe2+ in the blood plasma. Free
iron ions are toxic to most forms of life, and therefore transferrin binds Fe2+ and removes it from the blood.
There is an excess of transferrin present in the blood, and it has been shown that Ga3+ can also bind to this
glycoprotein but with a lower affinity than Fe3+. Once the binding capacity of gallium ions to transferrin is
exceeded, it is believed to circulate as gallate [Ga(OH)4−] [6a].

The therapeutic action of Ga3+ is very much based on the pharmacological activity of Fe3+ which it
mainly mimics. Ga3+ is transported via transferrin to areas of the body that require increased Fe3+ levels,
including proliferating cancer cells [11]. Ga3+ can interrupt the cell cycle and DNA synthesis by competing
with iron for the active sites in essential enzymes [12]. Ga3+ accumulates in the endosomes mediated by
transferrin uptake and transported into the cytosol, where it can bind to the enzyme ribonucleotide reductase.
Ribonucleotide reductase has been proposed as the main target for Ga3+. Binding to this enzyme will impair
DNA replication and ultimately lead to apoptosis [13]. In vitro studies have shown that Ga3+ can bind
directly to DNA [14].

4.4.4 Gallium nitrate – multivalent use

In clinical trials, gallium nitrate has proved to be highly active as an antitumour agent especially against
non-Hodgkin’s lymphoma and bladder cancer. The cytotoxic activity of gallium nitrate has been demonstrated
as single agent and as part of combination therapy, for example, together with fluorouracil. Gallium nitrate
shows a relatively low toxicity and does not produce myelosuppression, which is a significant advantage
over other traditional anticancer agents. Furthermore, it does not appear to show any cross-resistance with
conventional chemotherapeutic agents (Figure 4.8) [15].

These studies have also shown that gallium nitrate is able to decrease serum calcium levels in patients with
tumour-induced hypercalcaemia. Subsequently, several studies have been carried out comparing traditional
bisphosphonate drugs with gallium nitrate in their ability to decrease the calcium levels that are elevated as a
result of cancer. Based on the clinical efficacy, gallium nitrate injections (Ganite™) was granted approval by
the FDA for the treatment of cancer-associated hypercalcaemia. Gallium nitrate is also believed to inhibit the
bone turnover and therefore to decrease osteolysis, the active reabsorption of bone material, in patients with
bone metastasis secondary to other cancers.

O
Ga3+

N
O O3

Figure 4.8 Gallium nitrate

The Boron Group – Group 13 79

O N
O N

Ga

NO

Figure 4.9 Chemical structure of gallium 8-quinolate

4.4.5 Gallium 8-quinolinolate

Gallium 8-quinolinolate is a hexacoordinated Ga3+ complex in which the central gallium atom is coordinated
by three quinolinolate groups. It was developed as an orally available anticancer agent. It was successfully
tested in vitro against lung cancer and in transplanted rats against Walker carcinosarcoma [16]. Main side
effects were detected in experiments on mice at doses of 125 mg/kg/day. These included leukopaenia and
some fatalities. The highest concentrations Ga3+ were found in the bone, liver and spleen (Figure 4.9) [6a].

Preclinical studies have established the IC50 values for a single-agent activity in the lower micromolar
range for a variety of cancer cell lines. These cell lines include human lung adenocarcinoma, where gal-
lium 8-quinolinolate was shown to be 10 times more potent than gallium nitrate. Other cell lines include
melanoma and ovarian, colon and breast cancer. The inhibitory effect appears to be dose dependent and not
time dependent.

Gallium 8-quinolinolate entered phase I clinical trials under the drug name KP46 in 2004 in order to estab-
lish its safety and toxicity profile. KP46 was orally administered as a tablet, containing 10–30%w/w. Dose
up to 480 mg/m2 were given to patients with advanced solid malignant tumours. The drug was well tolerated
and preliminary success was seen in patients with renal cell cancer.

4.4.6 Gallium maltolate

Gallium maltolate [tris(3-hydroxy-2-methyl-4H-pyran-4-onato)gallium(III)] is a coordination complex con-
taining a central Ga3+ ion and three maltolate (deprotonated maltol) groups. Clinical studies have shown
that oral administration of gallium maltolate leads to significantly increased bioavailability compared to gal-
lium chloride. The oral bioavailability is estimated to 25–57% in comparison to 2% for gallium chloride
(Figure 4.10) [17].

Phase I clinical trials on healthy humans showed that doses were well tolerated up to 500 mg. Further-
more, the results suggested the possibility of a once-per-day treatment option as a result of the half-life
of the drug in the blood plasma (17–21 h). Orally administered gallium maltolate is excreted significantly
more slowly via the kidneys than gallium nitrate injected intravenously. It has been proposed that rapid intra-
venous administration leads to the formation of gallate, which is quickly cleared as a small molecule via the
kidneys. In contrast, oral administration leads to a slow loading of the blood plasma, and Ga3+ is bound to

80 Essentials of Inorganic Chemistry O
O
O
OO

Ga
OO

O

O

Figure 4.10 Chemical structure of gallium maltolate

transferrin. This may lead to a different mechanism of excretion, leading to a reduction in renal toxicity. Also,
the transferrin-bound Ga3+ has the potential to be directly transported to the cancer cell without causing sig-
nificant side effects. Therefore, an oral administration seems to be superior to parenteral administration [17].

4.4.7 Toxicity and administration
Gallium nitrate is usually administered as a continuous intravenous infusion (200 mg/m2 for 5 days) for the
treatment of cancer-induced hypercalcaemia. This dose is well tolerated even by elderly patients. Higher doses
are usually used in the treatment of cancer. Renal toxicities being the dose-limiting factor are normally seen
when gallium nitrate is administered as a brief intravenous infusion. With the long-term regime as described
above, diarrhoea is the most common side effect. Renal toxicity can normally be minimised by adequate
hydration of the patient [15].

The advantage of gallium nitrate therapies is that platelet count and white blood cell counts are not sup-
pressed, which means that no myelosuppression takes place, which represents a major advantage over con-
ventional chemotherapeutic agents [15].

The Boron Group – Group 13 81

4.5 Exercises

4.5.1 Draw the Lewis structure or chemical formula of the following aluminium-based drugs

(a) Aluminium acetate
(b) Aluminium chloride
(c) Aluminium oxide

4.5.2 Research different antacids mixtures, state their content and calculate the weight/volume per-
centage (%w/v) for each active pharmaceutical ingredient (API).

Drug name Aluminium %w/v Magnesium %w/v

hydroxide hydroxide

Maalox® 220 mg/5 ml ??? 195 mg/5 ml ???

4.5.3 A typical antacid capsule contains 475 mg aluminium hydroxide as the active ingredient.
How many milligrams of stomach acid (HCl) can be neutralised by one tablet?

4.5.4 An aluminium hydroxide suspension (30 ml) containing 500 mg/5 ml aluminium hydroxide is
prescribed to the patient. The prescription states the patient has to take 30 ml four times a day.

(a) What is the chemical formula of aluminium hydroxide?
(b) How many grams of Al3+ is given to the patient per single dose?
(c) What is the weight/volume percentage (%w/v) of aluminium hydroxide in the suspension?

4.5.5 Write the chemical equations explaining the amphoteric behaviour of gallium hydroxide.